About Chemistry
Upcoming SlideShare
Loading in...5
×

Like this? Share it with your network

Share

About Chemistry

  • 915 views
Uploaded on

This file contain the notes for Class 9 students, Shorts/Details Questions Answers

This file contain the notes for Class 9 students, Shorts/Details Questions Answers

More in: Education , Technology
  • Full Name Full Name Comment goes here.
    Are you sure you want to
    Your message goes here
No Downloads

Views

Total Views
915
On Slideshare
914
From Embeds
1
Number of Embeds
1

Actions

Shares
Downloads
13
Comments
1
Likes
1

Embeds 1

http://www.pinterest.com 1

Report content

Flagged as inappropriate Flag as inappropriate
Flag as inappropriate

Select your reason for flagging this presentation as inappropriate.

Cancel
    No notes for slide

Transcript

  • 1. 1 Q - 1) Define chemistry? Write the name of its branches: Ans: Chemistry is the branch of science that deals with the properties, compositions and structure of matter. It also deals with principles and the laws governing the changes involved in the matter. The names of its branches are as follow: 1. Physical chemistry 2. Organic chemistry 3. Inorganic chemistry 4. Bio chemistry 5. Analytical chemistry 6. Industrial or applies chemistry 7. Nuclear chemistry 8. Polymeric chemistry 9. Environmental chemistry Describe the branches of chemistry: Physical chemistry: It deals with the laws and principles governing the combination of atoms and molecules in chemical reactions. Organic chemistry: It is the study of hydrocarbons and their derivatives. It also deals with carbon compounds with the exception of CO2, CO etc. Inorganic chemistry: It is the study of elements and their compounds generally obtained from non – living organisms i.e. minerals. Biochemistry: It is the study of compounds of living organism i.e. plants, animals and their metabolism in the living body. Analytical chemistry: It is the study of methods and techniques involved to determine kind, quality and quantity of various components in given substance. Industrial or Applied chemistry: Itisthestudyofdifferentchemicalprocesses involved in the chemical industries for the manufacture of synthetic products like glass, cement, paper, soda ash, fertilizers, medicines etc. Nuclear chemistry: It is the study of changes occurring in the nuclei of atoms, accompanied by the emission of invisible radiations. Polymeric chemistry: It is the studyof polymerization and the products obtained through the process of polymerization such as plastic, synthetic fibers, papers etc. Environmental chemistry: It is the study of interaction of chemical materials and their effects on the environment of animals and plants. Q - 2) What is scientific method? Define Hypothesis and Theory. A particular methodthat is usedto searchfor facts andfigures is called Scientific Method. It consists of observation, hypothesis, theory and scientific laws. It may be possible answer of the problems after observation. It is also a trail idea for accumulating more knowledge or facts. It is the verified result of hypothesis, which is obtained with the help of careful experimentation. chemistry INTRODUCTION TO CHEMISTRY Chapter IX One
  • 2. 2 Q - 3) What is Scientific method? Write the name of its four stages and describe them: Scientific Method: A specific method that is used to search for truth or facts. And: It is based on observation, hypothesis, theory and scientific laws. Observation: It is the basic tool for explaining a phenomenon but it may vary from person to person according to his skill. Hypothesis: The statement may be possible answer of the problems after observation. It is also the trail idea for accumulating of more knowledge or facts. Theory: It is the verified results by hypothesis, which is obtained with the help of careful experimentation. Scientific law: When a theory is tasted repeatedly and is found to fit according to facts, gave valued predictions is called Scientific law. Q - 4) Write down the contribution of Muslim scientists in the field of Chemistry: Jabir – ibne – Hayyan: (721 – 803 A.D.) He is known as the father of chemistry.  He invented experimental methods for preparation of Nitric acid, Hydrochloric acid and white lead.  He developed method for extraction of metal from their ores.  He developed method for dyeing clothes. Al – Razi: (862 – 930 A.D.) He was physician, surgeon, philosopher & alchemist.  He was the first person who used opium as an anesthesia.  He prepared ethyl alcohol by fermentation process.  He divided the substances into living and non- living origins. Al – Beruni: (973 – 1048 A.D.) He contributed in mathematics, physics, chemistry, metaphysics, geography and history.  He determined the densities of different substances. Ibne – Sina: (980 – 1037A.D.) He contributed in the field of mathematics, medicine, medical chemistry, philosophy and astronomy. Q - 5) What is scientific law? Scientific Law: When a theory is tasted repeatedly and found to fit according to facts, and gave valued predictions is called Scientific Law.
  • 3. 3 Q - 6) Write down the contributions of Modern scientists in the field of Chemistry: Robert Boyle: (1627 – 1961 A.D.) He is known as father of modern chemistry.  He was first to put forward the idea that chemistry should be regarded as a systematic investigation of nature promoting knowledge.  He tried to purify chemicals to obtain reproducible reactions. J. Priestly: (1733 – 1804 A.D.)  He discovered oxygen, Sulphur dioxide and hydrogen chloride. J. J. Berzelius: (1779 – 1848 A.D.)  He introduced the idea of symbols, formulae and chemical equation to make the study more systematic. Arrhenius: (1859 – 1927 A.D.) and M. Faraday: (1791 – 1867A.D.)  They both gave the Ionic theory and Laws of electrolysis. Scheele: (1742 – 1786 A.D.)  He discovered chlorine. J. Black: (1728 – 1799 A.D.)  He made study of carbon dioxide Cavendish: (1731 – 1810 A.D.)  He discovered Hydrogen. John Dalton: (1766 – 1844 A.D.)  He gave atomic mass theory. Mendeleev: (1824 – 1907 A.D.)  He discovered the periodic arrangements of elements. Lavoisier: (1743 – 1794 A.D.)  He discovered that oxygen constituted about one-fifth of air. Q - 7) What important role chemistry plays in the society: Ans: Chemistry plays very important role in the society to prepare such as food, synthetic fibers, plastics, medicines, soap, detergents, cosmetics, cement, fertilizers, glass etc.    Q - 8) Write three significant reasons to study chemistry? It has important practical applications in the society the development of life saving drugs is one and complete list would touch upon most areas of modern technology. It is an intellectual enterprise, a way of explaining our material world. It figures prominently in other field, such as in biology in the advancement of medicines and useful intellectual tool for making important decisions
  • 4. 4 Q - 1) Sate the following laws: “It states that in any chemical reaction the initial weight of reacting substances is equal to the final weight of the products.” “It states that the different samples of the same compound always contain the same elements combined together in the same proportions by mass.” “It states that if two elements combine to form more than one compound. The masses of one element that combine with a fixed mass of the other element are in the ratio of small whole numbers or simple multiple ratios.” “It states that the two different elements separately combine with the fixed mass of third element, the proportions shall be either in the same ratio or simple multiple of it.” Q - 2) State law of conservation of mass with Landolt’s experiments Law of conservation of Mass: “It states that in any chemical reaction the initial weight of reacting substances is equal to the final weight of the products.” Landolt’s experiment: German Chemist H. Landolt’s tested different chemicals to verify Law of conservation of mass. 1. He took H – shaped tube that has two limbs. 2. He filled limbs A and B with Silver Nitrate (AgNO3) & Hydrochloric acid (HCl) respectively. 3. The tube was sealed and weighted initially. 4. The reactants were mixed by inverting and shaking the tube. 5. After mixing the tube weighted again and he observed that weight, remain same.  Chemical Reaction: AgNO3 + HCl → AgCl + HNO3 Q - 3) State the law of Constant, Definite proportion with example? Law of Definite proportions: “It states that the different samples of the same compound always contain the same elements combined together in the same proportions by mass. Example: Every sample of pure water, through prepared in laboratory or obtained from rain, river, or water pump contains one part of Hydrogen (H) and eight parts of Oxygen (O). Like H2O = 2 : 16 or 1 : 8 chemistry CHEMICAL COMBINATION Chapter IX TWO
  • 5. 5 Q - 4) State the law of multiple proportions with example: Law of Multiple proportions: “It states that if two elements combine to form more than one compound. The masses of one element that combine with a fixed mass of the other element are in the ratio of small whole numbers or simple multiple ratios.” Example: Carbon (C) form two stable compounds with Oxygen (O) namely carbon monoxide (CO) and carbon dioxide CO2. Compound Mass of Carbon (C) Mass of Oxygen(0) Ratio of Oxygen (0) CO 12 16 1 CO2 12 32 2 Q - 5) State the law of Reciprocal proportion with example? “It states that the two different elements separately combine with the fixed mass of third element, the proportions shall be either in the same ratio or simple multiple of it.” Example: When two elements Carbon (C) and Oxygen (O) separately combine with fixed mass of Hydrogen (H) to form methane CH4 and water H4O.  Ratio of H is 1 in CH4 & H2O  Ratio of C is 3 in CH4 & CO2  Ratio of O is 8 in H2O & CO2 H CH4 H2O 12 : 4 2 : 16 3 : 1 CO2 1 : 8 C O12 : 32 3 : 8 Q - 6) What is Einstein theory? & Write reaction B/W mass & energy: Einstein Theory: The famous physicist and mathematician Albert Einstein proposed the relation between mass and energy. i.e. “There is no detectable gain or loss of mass in a chemical reaction”. Mathematically E = mc2 where “E” is energy, “m” is mass and “c” is speed of light. Q - 7) What is chemical reaction? Define and give example each of the following: Chemical Reaction: It is a change in which the composition of substances is changed and new substances formed. It has five types as follows. (1) Decomposition Reaction: It is the reaction in that breaks down a compound or chemical substances to form two or more substances. Heat CaCO3 CaO + CO2
  • 6. 6 (2) Addition Reaction: It is the reaction in which two or more substances combine together to form new substances. Example: CaO + CO2 CaCO3 Example: Na + Cl2 2NaCl (3) Single Displacement Reaction: It is the reaction in which one atom or a group of atoms of a compound replaced by another atom or group of atoms. Example: 2Na + 2H2O 2NaOH + H2 Example: Zn + 2HCl ZnCl2 + H2 (4) Double Displacement Reaction: It is the reaction in that two compounds exchange their partners to form new compounds. Example: NaCl + AgNO3 NaNO3 + AgCl Example: CaCl2 + Na2CO3 2NaCl + CaCO3 (5) Combustion Reaction: It is the reaction in that substances react with either free oxygen or oxygen of air, with rapid release of heat and flame. Example: CH4 + O2 CO2 + 2H2 + ∆H (Heat) Example: C + O2 CO2 + ∆H (Heat) Q - 8) What is chemical reaction? Define & give one example each: (1) Addition reaction (2) Double Displacement reaction Chemical Reaction: It is a change in which the composition of substances changes to forms new substances. (1) Addition Reaction: It is the reaction in which two or more substances combine together to form new substances. Example: CaO + CO2 CaCO3 (2) Double Displacement Reaction:It is the reaction in which, two compounds exchange their partners to form new compounds. Example: NaCl + AgNO3 NaNO3 + AgCl Q - 9) What is chemical reaction? Define & give one example of each: (1) Combustion reaction (2) Single Displacement reaction Chemical Reaction: It is a change in which the composition of substances is changed and new substances formed. (1) Combustion Reaction: It is the reaction in which, substances react with either free oxygen or oxygen of air, with rapid release of heat and flame. Example: CH4 + O2 CO2 + 2H2 + ∆H (Heat) (2) Single Displacement Reaction: It is the reaction in which one atom or a group of atoms of a compound replaced by another atom or group of atoms. Example: 2Na + 2H2O 2NaOH + H2
  • 7. 7 Q - 10) What is Empirical formula? Give an example: Empirical Formula: (E.F) The formula that describes the smallest or least ratio of the combining atoms of different elements presents in a molecule. Example:The empirical formula of Benzene is CH Mathematically Example:The empirical formula of Glucose is CH2O 𝑬. 𝑭 = 𝑴. 𝑭. 𝒏 Q - 11) What is Molecular formula? Give an example: Molecular Formula: (M.F) The formula that describes actual number and type of the combining atoms of all elements presents in a molecule. Example: The molecular formula of Benzene is C6H6 Mathematically Example: The molecular formula of Glucose is C6H12O6 𝑴. 𝑭 = ( 𝑬. 𝑭) 𝒏 Q - 12) Define the following: (i) Atomic mass (ii) Mole (iii) Molar Mass (iv) Avogadro’s Number (i) Atomic mass: It is the average mass of naturally occurring isotopes that is compared to the mass of one atom of carbon – 12 a.m.u. (ii) Mole: It is amount of substances contains as many elementary particles like atoms, ions, molecules etc. It is denoted by n. Formula n = m M (iii) Molar mass: It is the mass in grams of one mole of substances. Such as (molecular mass, formula mas, atomic mass) etc. (iv) Avogadro’s number: It is a constant number that is equal to 6.02 × 1023 of one mole of any substances. It is denoted by NA. Q - 13) Can one substance have the same empirical formula and molecular formula? Explain with examples: Ans: Yes, one substance can have the same empirical formula and molecular formula because they have the simplest and least whole number and have no difference between their molecular and empirical formulas. Example: Name Empirical Formula Molecular Formula Sodium Hydroxide NaOH NaOH Sugar C12H22O11 C12H22O11 Methane CH4 CH4 Potassium Nitrate KNO3 KNO3
  • 8. 8 Q - 14) What is chemical equation? What is a co-efficient and expression? Give an example of balance equation: Chemical Equation: It is short hand method of describing the chemical reaction, in terms of symbols and formulae of the substances involved in a chemical reaction. Co – efficient: The number, in front of the formulas in a chemical equation is called Co – efficient. It is present before molecule or atoms. Expression: the latters, (g), (l) and (s) are placed as subscript indicates the states of products and reactants. Balanced Equation: Example: MnO2 2KClO3 (S) 2KCl(s) + 3O2(g) Q - 15) The value of carbon in periodic table is 12.011 a.m.u rather than 12.00 a.m.u. Explain it: Ans: The value of carbon in periodic table is 12.011 a.m.u rather than 12.00 a.m.u because naturally Carbon has two isotopes C– 12 with 98.889% and C-13 with 1.111. Thus the atomic mass of Carbon (C) atoms becomes 12.011 a.m.u. Average/Sum of Carbon isotopes = 12×98.889+13×1.111 100 = 12.011a.m.u. Q - 16) Differentiate Empirical formula and molecular formula: Empirical formula Molecular formula It describes smallest or least ratio of the elements in compound. It describes actual number and type of elements in compound Some compound may have same empirical formula. No, any two compounds may have the same molecular formula. Example: Empirical formula of Acetic acid and glucose is same as CH2O Example: Molecular formula of Acetic acid is CH3COOH and Glucose is C6H12O6 Molecular mass Formula mass The term molecular mass cannot be used with ionic compounds because there are no discrete molecule in ionic compounds. The term formula mass can be used with either molecular compounds or ionic compounds. It is the sum of atomic masses of all atoms in a molecular formula of substances. It is the sum of atomic masses of all atoms in a formula unit of substances. Molecular mass of CO2 is 44 a.m.u. Formula mass of NaCl is 58.5 a.m.u.
  • 9. 9 Q - 1) Write down the main postulates of Dolton Atomic theory: Dolton Atomic Theory:  All elements are made up of small invisible, indestructible particles are called atoms.  All atoms of a given element are identical in all respects having same size, mass and chemical properties. However, the atoms of one element differ from the atoms of other element.  Compounds are formed when atoms of more than one element combine in a simple whole number ratio.  A chemical reaction is a rearrangement of atoms but atoms themselves are not changed this means that atoms are neither created nor destroyed in chemical reactions. Q - 2) Describe Modern Atomic theory: Modern Atomic Theory: According to modern atomic theory, Atom is a complex organization and composed of even smaller particles called sub – atomic particles (Fundamentals particles). They are Electrons, Protons and Neutrons. Q - 3) Write down the properties of Cathode Rays? Properties of Cathode Rays: i. They cast shadows of objects towards anode and travels in straight lines. ii. They cause a light paddle wheel to rotate. Showing that they are material particles. iii. They are negatively charged particles and deflected towards positive plate. iv. The (Charge/Mass) e/m ratio of them is 1.7588 x 108 c/g. v. They possess Kinetic Energy so they can produce mechanical pressure. vi. They are invisible cause some material to glow or produce fluorescence. Q - 4) Write down the properties of Positive Rays? Properties of Positive Rays: i. They also travel in straight line from anode to cathode. ii. They cause a light paddle wheel to rotate. Showing that they are material particles. iii. They are positively charged particles and deflected towards negative plate. iv. The (Charge/Mass) e/m ratio of them is much smaller than electrons. v. Their e/m ratio varies with the nature of gas in the tube. vi. They have +1.6022 x 10–19 coulomb charge. CHEMISTRY ATOMIC STRUCTURE CHAPTER IX THREE
  • 10. 10 Q - 5) Describe the discovery of electrons (Cathode Rays) Discovery of Electrons: The British physicist J.J. Thomson discovered electrons by an experiment in 1897. For that he used a glass tube (Discharge tube) fitted with two metal electrodes connected to a high voltage source and a vacuum pump. He evacuated the tube and passed a current of high potential between the electrodes at very low pressure.  He observed that streaks of bluish light extending from negative electrode (cathode) towards positive electrode (anode).  These rays travelled in straight line and glow at the opposite end where they stroked.  He showed that these rays were deflected towards the positive plate in electric and magnetic field.  He named the rays “Electron” and they have negative (–) charge. Q - 6) Describe the discovery of Protons: (Positive Rays) what are protons and how were these produced? Discovery of Protons: The German physicist Goldstein discovered protons by an experiment in 1886. For that, he used a glass tube (Discharge tube) with perforated cathode, fitted two metallic electrodes connected to a high voltage source. He conducted a series of process during his experiment. He perforated cathode containing a gas at low pressure and passed current with high potential source. He observed that rays containing of positively charged particles traveled from anode to cathode. These rays also travelled in straight line towards negative plate.  He showed that positively charged particles was equal to negatively charged particles.  He named these rays “Protons” and they have positive (+) charge.
  • 11. 11 Q - 7) Describe the discovery of Neutrons: Discovery of Neutrons: The English physicist James Chadwick discovered Neutrons through artificial radioactivity in 1932. For that he bombarded Be (Beryllium) with alpha particles, neutrons were appeared. Q - 8) Describes the properties of electron, proton and neutron: Electron:  Electron is negatively charged particle.  It has electric charge about – 1.602 x 10 – 19 coulombs.  It moves around the nucleus.  It has 9.109 x 10 –31 kg mass. i.e. 1 1836 Of protons. Proton:  Proton is positively charged particle.  It has electric charge about + 1.602 x 10– 19 coulombs.  It lies in the nucleus.  It has 1.672 x 10 –27mass. i.e. 1836 times of electron. Neutron:  Neutral particle has no charge.  It lies in the nucleus.  It has 1.762 x 10 –27mass. i.e. 1836 times of electron. Q - 9) What are the names of sub – atomic particles? What are the masses in atomic mass units (a.m.u) and what is unit charges each? Ans: The names of three Subs – atomic particles are electron, proton and neutron. Electron: Its mass is 0.0005485 a.m.u and unit charge +1. Proton: Its mass is 1.007276 a.m.u and unit charge –1. Neutron: Its mass is 1.008664 a.m.u and unit charge 0 (None). Q - 10) Define (i) Isotopes (ii) Mass number (iii) Atomic number Isotopes: Atoms of the same element having the same atomic number but different atomic masses are called isotopes. Such as Protium H1 1 , Deuterium D1 2 and Tritium T1 3 are isotopes of Hydrogen (H). Mass Number: The total sum of the protons and neutrons in the nucleus of an atom is called the mass number and it is denoted by A. Atomic Number: The number of protons in the nucleus or number of electrons outside the nucleus of an atom is called atomic number and it is denoted by Z.
  • 12. 12 Q - 11) What is meant by Radioactivity? Draw a labeled diagram showing the separation of Alpha, Beta and Gamma rays: And also write the properties of each: Radioactivity: The process in which emission of fundamental particles, in the form of radiations takes place is called radioactivity. The British physics Ernest Rutherford in 1902 determined the nature of radioactive rays by an experiment and showed that it is composed of three types of rays. For that, he placed a sample of radioactive substances in a lead block, between the two oppositely charged plates. The radiations were resoled into three components.  First component was deflected towards negative plate, having (+) charge & name as 𝛼– rays.  Second component was deflected towards positive plate, having (–) charge & name as 𝛽– rays.  Third component was not deflected towards any plate, having no charge & name as 𝛾– rays. 𝜶 – i. 𝛼– Rays deflect towards negative plate, and having (+3.2 × 10−19 𝑐 ) charge. ii. They have heavy masses that are why they cannot pass through thick metal foil. iii. They are strong ionizing agents. iv. They are Helium nuclei. 𝜷 – i. 𝛽– Rays deflect towards positive plate, and having (−1.6 × 10−19 𝐶 ) charge. ii. They have great power to penetrate that is why they can also pass through thick metal foil. iii. They are weak ionizing agents. iv. They are electrons. 𝜸 – i. 𝛾 – Rays do not deflect towards any plate, and they are neutral. ii. They have greater power to penetrate that is why they can pass through thicker metal foil than 𝛽 – rays. iii. They are weaker ionizing agents than 𝛽 – rays. iv. They are electromagnetic radiations.
  • 13. 13 Q - 12) Describe the Rutherford’s experiment which led him to the discovery of the nucleus of an atom? Or Describes Rutherford’s experiment of Gold metal foil: Discovery of Nucleus: Rutherford discovered Nucleus of an atom by an experiment in 1911. For that, he used Gold metal foil and Alpha particles. He bombarded a beam of alpha 𝛼 – particles through a very thin gold metal foil. He observed that most of alpha 𝛼 – particles passed through it without any deflection. Some of them deflected at large angles and very few of them bounced back. He concluded two things Volume and presence of Nucleus.  The volume occupied by an atom must be largely empty as most of the 𝛼 – particles passed through the foil unbounded.  The positive charge, in the atom is concentrated in extremely dense region that he called the nucleus. Rutherford atomic model: Accordingtohim,anatomconsiststwoparts. (i) Nucleus (ii) Extra Nuclear part Nucleus:  It has positive (+) charge because proton and neutron reside in it.  The weight of an atom is concentrated in the nucleus. Extra nuclear part:  The electron is revolving around the nucleus in the extra nuclear part in various orbits, which are also called as shells or energy levels. Defects of atomic model:  Electrons are continuously revolving around the nucleus that may cause lose their energy and fall into nucleus.  If revolving electrons emit energy continuously, then there would be a nonstop spectrum but we get it from the atoms of elements.
  • 14. 14 Q - 13) Explain Neil Bohr’s Atomic theory: Neil Bohr’s atomic theory: The Danish physicist offered theoretical assumptions for atomic structure in 1913. The important assumptions are as following. i. Electrons in atoms move only in certain allowed energy levels (energy states) so they will not fall in the nucleus. ii. The atom radiates energy as light when the electron jumps higher energy level (E2) to lower energy level (E1). iii. The quantity of energy radiated is in discrete quantity, called Quanta. iv. A quantum of energy is directly proportional to the frequency of radiation. i.e. ∆𝐸 = 𝐸2 − 𝐸1 = ℎ𝑣 Where ℎ = Planks constant and 𝑣 = frequency of radiation. Q - 14) Write down the application of Isotopes: Application of Isotopes or Uses of isotopes: i. Isotopes are used as tracers in physical, chemical and biological researches. ii. Isotopes are also used in treatments and diagnoses of disease like cancer. Q - 15) C – 14 and N – 14 both have same mass number yet they are different elements, Explain: Ans: C – 14 and N – 14 both have same mass number yet they are different elements because of their different atomic numbers as atomic number of Carbon is 6 and Nitrogen is 7 and C – 14 is isotope of Carbon. Q - 16) In what way isotopes of a given element differ from each other? Ans: It is because that the different isotopes of the same elements differ only in the number of neutrons in the nucleus of an atom. Q - 17) What are isotopes? Write the symbol of Protium, deuterium and tritium with their atomic numbers and mass numbers. Isotopes: Atoms of the same element having the same atomic number but different atomic masses are called isotopes. Name of Isotopes Symbols Atomic Number Mass number Protium H1 1 H 1 1 Deuterium D1 2 D 1 2 Tritium T1 3 T 1 3
  • 15. 15 Q - 1) Define the following: (i) Doberneir’s rule of traid (ii) Modern periodic law (iii) Periodicity (iv) Electronegativity (v) Electron Affinity (vi) Atomic Radii (vii) Ionization Energy (viii) Period (ix) Group i. Doberneir’s Rule of traid: It states, “Central atom of each set of traid had an atomic mass almost equal to the arithmetical mean of the atomic masses of the two elements.” ii. Modern Periodic law: It states “The physical and chemical properties of elements are periodic functions of their atomic weight”. iii. Periodicity: The repetition of Physical and Chemical properties of elements periodically is called periodicity of elements. iv. Electronegativity: The relative power of an atom to attract the shared pair of electrons towards itself is called electronegativity. The most electronegative atom is Florien (F) with electronegativity = 4. v. Electron Affinity: It is defined as the energy change that occurs when an electron is gained by an atom on the gaseous state. vi. Atomic Radii: It is defined as half the distance between two adjacent nuclei of two similar atoms in touch with each other. It is measured in Angstrom unit Ao . vii. Ionization Energy: It is defined as the minimum energy required removing an electron from a gaseous atom in its ground state. It is measured in K.Jol/Mol or Electron volt per atom. Ev/atm. viii. Period: The modern periodic table contains seven horizontal rows called periods. The elements within a period have dissimilar properties from left to right. ix. Group: The modern periodic table contains sixteen vertical columns called groups. It is divided into two groups A and B. the elements of Group A are called main or representative elements. And elements of Group B are called Transition Elements. chemistry PERIODICITY OF ELEMENT Chapter IX Four
  • 16. 16 Q - 2) Write characteristics Metals, Non – metals and Metalloids: Metals: (i) They are electropositive and lose electrons to form cat – ions. (ii) They are basic oxides and good conductors of heat and electricity. (iii) All of them have luster and are malleable and ductile. Non – Metals: (i) They are electronegative and Gain electrons to form an – ions. (ii) They are acidic oxides and bad conductors of heat and electricity. (iii) Most of them are gases and have no definite shape and volume. Metalloids: (i) They have dual character and shows properties of both metals and non - metals. (ii) They have basis as well as acidic nature. (iii) They are semiconductors and brittle rather than malleable. Q - 3) Describe properties of Transition elements: (Group IB to VIIB) (i) They are metals and they have incomplete valence shall. (ii) In chemical reactions they show more than one valences. (iii) These elements in compounds having characteristic colors. (iv) They have ability to form complex ions by coordination. Q - 4) Describe the properties of Group I A, to Group VIII A: Group I – A: (Lithium Family / Alkali Metals) They are called alkali metals because they form water soluble base such as NaOH, KOH. Their atomic radii, atomic volumes, ionic radii increase from Li to Cs. Their melting and boiling point decrease downward. They are highly reactive metals with low melting points. This group contains Lithium (Li), Sodium (Na), Potassium (K), Rubidium (Rb), Cesium (Cs) and Francium (Fr). Their valance shell contains only one electron and they lose it in chemical reactions. Group II – A: (Beryllium Family / Alkaline Earth Metals) They are called alkaline earth metals and they are a bit harder, have higher melting and boiling point. Their valance shell contains only two electron and they lose them in chemical reactions. They have smaller atomic radii, atomic volumes than Alkali metals. Downward they do not show a regular trend in melting, boiling points and densities. This group contains Beryllium (Br), Magnesium (Mg), Calcium (Ca), Strontium (Sr), Barium (Ba) and Radium (Ra).
  • 17. 17 Group III – A: (Boron Family) They are highly reactive elements except Boron (Br), having metallic character. Their valance shell contains only 3 electrons. It means their valance is +3 They have moderate tendency to form compounds. Born is metalloid having properties of metal and non – metal. Downward they decrease their ionization potential and increase size of atoms. This groups contains Boron (B), Aluminium (Al), Gallium (Ga), Indium (In) and thallium (Tl) Group IV – A: (Carbon Family) They are highly reactive elements except Boron (Br), having metallic character. Their valance shell contains 4 electrons, except Tin (Sn) and Lead (Pb) exhibit a variable valance 2 and 4. The elements of this group are metal, non – metal and metalloid. Such as Sn and Pb are metals, C is non – metal and Si and Ge are metalloids. Down the group, their atomic radii and volumes increase due to addition of new shells. Carbon (C) and Tin (Sn) exist in different allotropic forms. This groups contains Carbon (C), Silicon (Si), Germanium (Ge), Tin (Sn) and Lead (Pb) Group V – A: (Nitrogen Family) Their valance shell contains 5 electrons. Nitrogen have tendency to except three electrons and forms number of oxides as NO, NO2, N2O4 etc. The elements of this group are metal, non – metal and metalloid. Such as Bi is metal, N and P are non – metals and As and Sb are metalloids. The elements of this group are quite reactive, having quite tendency to form compounds. Except Nitrogen (N) all exist in more than one allotropic form. This group contains Nitrogen (N), Phosphorus (P), Arsenic (As), Antimony (Sb) and Bismuth (Bi). Group VI – A: (Oxygen Family) Their valance shell contains 6 electrons. Oxygen (O) and Sulphur (S) form divalent negative ions O2- and S2- The elements of this group are metal, non – metal and metalloid. Such as (Po) is metal, O and S are Non – metals and Se and Te are metalloids. The elements of this group are quite reactive, having quite tendency to form compounds. All of them exhibit the property of allotropy. This group contains Oxygen (O), Sulphur (S), Selenium (Se), Tellurium (Te) and Polonium (Po).
  • 18. 18 Group VII – A: (The Halogens) Their valance shell contains 7 electrons. Oxygen (O) and Sulphur (S) form divalent negative ions O2- and S2- Except Astatine (At) which is metalloid, All other are non – metals and exist as diatomic molecules. They have high ionization energies and large negative electron affinities. They easily accept an electron to form halide ions i.e. F1- , Cl1- Br1- and I1- All of them exhibit the property of allotropy. Such as Oxygen (O2)and Ozone (O3) This group contains Fluorine (F), Chlorine (Cl), Bromine (Br), Iodine (I) and Astatine (At). Group VIII – A: (Inert and Noble Gases) Their valance shell contains 8 electrons. Except helium (He) which has two electrons. They are mono atomic and exist in gaseous state. WithexceptionofKrypton(Kr) andXenon(Xe)areslightlyreactiveunderdrasticcondition and rest all of them are very inert chemically. They are mostly chemically non – reactive because of great stability. Thisgroupcontains Helium(He),Neon(Ne),Argon(Ar),Krypton(Kr),Xenon(Xe)andRadon (Rn). Q - 5) If an element contains two shells only and its outer most shell contains five electrons then to which group the element belongs in the periodic table. Name the element. Predict its period. Ans: If an element contains two shells only and its outer most shell contains five electrons then the element belongs to group VA in the periodic table. The name of element is Nitrogen (N) and its period is 2. Q - 6) What do you mean by long form of periodic table? Explain some of its application: Ans: The modern periodic table is known as the long form of periodic table because elements are arranged in order of increasing atomic number. It contains seven horizontal rows called periods and sixteen vertical columns called groups. Application of periodic table: Prediction of new elements and their chemical characteristics is now possible. The classification of elements into periods and groups simplifies the study of chemistry. Suggestion for further research becomes available now. Reactivity of elements can be visualized.
  • 19. 19 Q - 7) Discuss some of the physical properties of the elements which exhibit periodicity. Ans: The physical properties of elements depend upon boiling point, melting point, densities, electric conductivity and hardness. Physical Properties of elements:  The elements of group IA, IIA, IIIA, IVA, VA, and VIA are metals and conduct electricity.  The atomic radii increase down the group due to addition of new shell in each atom.  Melting and boiling points in group IA decrease gradually from top to bottom  The elements are at right side Non – metals, at left side are metals and in the middle are metalloids. Q - 8) How does the modern periodic law differ from Mendeleev’s periodic law? Explain groups and periods in modern periodic table: Ans: According to The Modern periodic law “The physical and chemical properties of all elements are periodic functions of their “Atomic numbers” while By Mendeleev’s periodic law “The physical and chemical properties of elements are periodic function of their Atomic weights”. Groups: The modern periodic table contains sixteen vertical columns called groups. It is divided into two groups A and B. Elements of group “A” are called main or representative elements. And group B’s elements are called Transition Elements. Period: The modern periodic table contains seven horizontal rows called periods. The elements within a period have dissimilar properties from left to right. Q - 9) What do you understand by periodic classification of elements? What are the merits and demerits of classification of elements: Ans: According to the periodic classification of elements are arranged in the order of increasing atomic number in period, such that all elements having same number of valence electrons grow under the same vertical column called group. This also ensures periodicity in electronic configuration and in chemical properties.  The classification of element is based on fundamental property of elements.  It explains variations and similarities in the properties of elements in term of electronic configuration.  It provides clear demarcation of different kinds of elements such as active metals, metalloids, non – metals, transition elements, inert gases etc.  It relates position of an element to its electronic configuration in the valence shell.
  • 20. 20  The classification of elements is unable to reflect electronic configuration of many elements in transition group, in lanthanides and actinides.  It fails to accommodate lanthanides and actinides in the main body of the table.  Position of Hydrogen is unresolved. In addition, position of Helium amongst p – block elements is not fully justified. Q - 10) What do you understand by representative and transition elements? or What do mean by transition elements? Write the general properties/characteristics of transition elements. Representative elements: The elements of group “A” are called “Main elements” or “Representative elements” as the properties of these elements are represented by valence electrons. Transition elements: The elements of group “B” are called “Transition elements” because the properties of these elements show a gradual change or transition between the two sets of representative elements. 1. These elements have incomplete inner electron shells. 2. They have an ability to form complex ions by coordination. 3. They have strong interred atomic bonds and having characteristics color. 4. They are all metal and having variable valence with similar behavior. 5. In these elements, penultimate shell is also incomplete. Q - 11) What are Lanthanides and Actinides? Are they d or f – block elements: Lanthanides: The Fourteen inner elements of sixth period are called Lanthanides. They start from Ce to Lu. They have six electronic shells. Actinides: The Fourteen inner elements of seventh period are called Actinides. They start from Actinium Ac to Lr. They have seven electronic shells. Q - 12) State Mendeleev’s periodic law. Describe Mendeleev’s periodic table and also write down the merits and demerit of Mendeleev’s periodic table: Mendeleev’s Periodic Law: Statement: “The physical and chemical properties of elements are a periodic function of their atomic weights.”
  • 21. 21 SALIENT FEATURES OF TABLE:  It has eight vertical columns called groups and twelve horizontal rows called period.  Elements in each vertical column have similar properties.  Vacant spaces were left for the elements not discovered until then.  He proposed their names as eka – boron, eka – aluminum and eka – silicon.  The group number indicates the highest valence that can be attained by elements of that group.  It helps in systematic study of elements and forcefully proved the concept of periodicity.  The prediction of new elements was made possible.  It helped in correcting many doubtful atomic masses.  In his table, there was no place for isotopes of elements.  It fails to give idea of atomic structure.  Dissimilar elements were placed in same group i.e. Alkali metals (Li, Na, K, Rb, Cs, Fr) were placed with coinage metals (Ag, Cu, Au)  Similar elements placed in different groups. For example, Barium (Ba) and Lead (Pb) resemble in many properties but they are placed in separate groups. Q - 13) Explain Newland’s Law of octave. How this law provide the larger scope for the classification of the elements: Newland’s Law of octave: Statement: “If elements are arranged in the order of increasing atomic masses the eight element starting from a give one, has similar properties as first one i.e. its properties is a kind of repetition of the first, like the eight note on an octave of music” Larger Scope for Classification: Ans: This law provides the larger scope for the classification of elements by the arrangement of its elements with their physical and chemical properties at a regular interval that show the existence of periodicity of the Modern periodic table. Defects of Newland’s Law:  “Hydrogen” was not included in his sequence.  His sequence was failed because it worked for first sixteen elements.
  • 22. 22 Q - 1) Define Chemical bonding and Discuss how atoms unite and change into molecules: ChemicalBonding: The Force, which holds atoms together in a molecule or a crystal, is called a chemical bond. When one electron or more than one electron of any element combines with other or same group of elements, they change into molecules or compound. Example: In combination of NaCl (Sodium Chloride) an atom of Sodium (Na) transfers one outer most shell electron, becomes positive sodium ion (Na+) and an atom of chlorine gains that one electron to complete its octet, and becomes chloride negative ion (Cl– ). Na(2,8,1) Na+ (2,8) + e– Cl(2,8,7) + e– Cl – (2,8,8) Na+ + Cl– NaCl Q - 2) What are the main types of bond? Describe them: There is two Main type of bond (i) Ionic or electrovalent bond (ii) Covalent bond Ionic/Electrovalent Bond: In these types of bond, there is a complete transfer of one or more electrons from one atom to another. The atom that transfers electrons gets positive charge and the atom that gains electron gets negative charge. They also called ionic bond because they produce ion and conduct electricity when dissolve in water or melted. Covalent Bond: In these types of bond, each atom has to contribute equal number of unpaired electrons. The shared pair of electrons, which links the atoms in a molecule, is known as covalent bond. In covalent bond, the shared pair of electron is commonly expressed by single short line (—). Covalent bond has three types (i) Single covalent bond (ii) Double covalent bond and (iii) Triple covalent bond CHEMISTRY CHEMICAL BONDING CHAPTER IX FIVE
  • 23. 23 Q - 3) Define: (i) Single covalent bond (ii) Double covalent bond and (iii) Triple covalent bond? or What are the types of chemical bonding? Single CovalentBond: In this bond, only one pair of electrons is shared by the bonded atoms, in which each atom has to share one electron. This type of bond represents by single short line (–). For example: Formation of hydrogen H2 Double Covalent Bond: In this bond, only two pairs of electrons are shared by the bonded atoms, in which each atom has to share two unpaired electrons. This type of bond represents by double short lines (=). Formation of Double covalent bonds in O2 Triple Covalent Bond: In this bond, only three pairs of electrons are shared between the bonded atoms, in which each atom has to share three unpaid electrons. This type of bond represents by single short line (≡). Q - 4) What happens to electrons, when elements combine? Ans: When elements combines the valence electrons are either transferred from the outer shell of one atom to the outer shell of another atom or shared between them and it produces Chemical Bond. Q - 5) What part of the atom is involved in the formation of chemical bond? Ans: The outer most part (The valence electrons) of the atoms is involved in the formation of chemical bond. Q - 6) Explain with examples? How elements are united by electrovalent bond? Ans: The elements are unite by electro – Valente bond are one, two or three and the formation of their bonding is known as single, double or triple covalent bonds respectively. Example: HCl (single covalent bond), CO2 (double covalent bond) and C2H2 (triple covalent bond)
  • 24. 24 Q - 7) What common properties are shown by ionic compound? Ans: The elements are unite by electrovalent bond one, two or three and the formation of their bonding is known as single, double or triple covalent bonds respectively.  They held very strong electrostatic forces and they are solids at room temperature.  They have high melting and boiling points.  They are hard and brittle but easily broken.  In solid state, they do not conduct electricity but in liquid state they conduct it.  They are soluble in water (polar) and insoluble in non – polar solvents.  They are non – volatile in nature and generally inorganic compounds. Q - 8) What is meant by covalent bond? Write electronic formulas of any covalent molecules? COVALENTBOND: In these types of bond each atom has to contribute equal number of unpaired electrons. The shared pair of electrons, which links the atoms in a molecule, is known as covalent bond.  The formation of molecules of CO2 and C2H2. × C × × × × × × C HH × H C C H
  • 25. 25 Q - 9) Account for the fact that some covalent bonds are polar and while other is non – polar? Condition for Non – Polar Covalent Bond: According to the scale of Linus Pauling, if the difference in the electronegativity of bonded atoms is zero, then the bond is pure Covalent bond or Non – Polar bond. Examples: H – H, O = O, N ≡ N, H3C – CH3, H2C = CH2 and Cl – N – Cl | Cl Condition for Polar Covalent Bond: According to the scale of Linus Pauling, if the difference in the electronegativity of bonded atoms is up to 1.7, then the bond is called Polar covalent bond or partially ionic in character. Examples: H+𝛿 – Cl−𝛿 H+δ – Cl−δ – H−δ | H H | H – C+δ – Br−δ | H H+δ – O−δ | H+δ Q - 10) What is co – ordinate covalent bond? Explain with examples: Co – Covalent Bond: The Co – ordinate covalent bond is formed only when an atom with an unshared pair of electrons in its valence shell donates a pair of electrons to anther atom or ion that needs a pair of electrons to acquire a stable electronic configuration. Examples: The formation of hydronium ion (H3O+) from water molecule and hydrogen ion (H+) in which the oxygen atom of water acts as donor and hydrogen (H+) ion as accepter. Donor Accepter Co – ordinate covalent bond ×× H+ ×× H – O → H | H Formula H ∙ × o ×× H3O+ × H
  • 26. 26 Q - 11) Define the term covalent bond? How does a covalent bond differ from co – ordinate covalent bond? COVALENTBOND: In this type of bond, each atom has to contribute equal number of unpaired electrons. The shared pair of electrons, which links the atoms in a molecule, is known as covalent bond. CO – CONALENT BOND: The Co – ordinate covalent bond is formed only when an atom with an unshared pair of electrons in its valence shell donates a pair of electrons to anther atom or ion that needs a pair of electrons to acquire a stable electronic configuration. Covalent bond Co – ordinate covalent bond 1. The mutual sharing of electrons between atoms forms this. 1. This bond is formed by the one sided sharing of electrons. 2. This bond is formed between the similar or dissimilar atoms. 2. This bond formed between two unlike atoms. 3. The reaction of this bond is denoted by short line (– ), (=) & (≡). 3. An arrow sign denotes the reaction of this bond (⟶). 4. They may be polar or non – polar. 4. These bonds are always polar. 5. They are usually insoluble in water 5. They are sparingly soluble in water. Q - 12) Explain electronegativity with the help of Linus Pauling table: Electronegativity: The combining power of an atom to attract the shared pair of electrons towards itself is known as electronegativity. Explanation: Linus Pauling proposed the table of electronegativity of different elements, in which he gave fluorine (Fl) as an arbitrary standard value of electronegativity as 4 because it is most electronegative element. The electronegativity value of other elements is compared with fluorine. Values of E.N. of different Elements Li = 1.0 Be = 1.5 B = 2.0 C =2.5 N = 3.0 O = 3.5 F = 4.0 Na = 0.9 Mg = 1.2 Al = 1.5 Si = 1.8 P = 2.1 S = 2.5 Cl = 3.0 K =0.8 Ca = 1.0 Sc = 1.3 Ti = 1.5 V = 1.6 Cr = 1.6 Mn = 1.5 Fe = 1.8 Co = 1.8 Ni = 1.8 Cu = 1.9 Zn = 1.6 Ga = 1.6
  • 27. 27 Q - 13) Give characteristics of covalent compounds:  They are usually made up of discrete units (molecules) with a weak inter molecular forces.  In solid state there are weak Vander wall forces between the molecules.  They are often gases, liquids or soft solids with low melting points.  They usually have low melting and boiling points.  They are insoluble in water (polar) but soluble in organic solvents like benzene, ether, carbon tetra chloride etc. Q - 14) What is chemical bond? Define Ionic Bond. Explain the mechanism of the ionic bond in NaCl (Sodium Chloride)? Chemical Bond: It is the holding of group of atoms together to form molecules or solids. In addition, it is occurs when a group of atoms can lower its total energy by combining. Ionic or Electrovalent Bond: In this types of bond there is a complete transfer of one or more electrons from one atom to another. Examples: The formation of sodium chloride (NaCl) an atom of (Na) transfers one outer most shell electron, becomes positive sodium ion (Na+) and an atom of chlorine gains that one electron to complete its octet, and becomes chloride negative ion (Cl–). + – Na Cl Na+ Cl– Na(2,8,1) Na+ (2,8) + e– Cl(2,8,7) + e– Cl – (2,8,8) Na+ + Cl– Na Cl or Na+ Cl–
  • 28. 28 Q - 15) What is coordinate Covalent Bond? How a coordinate covalent bond is formed between ammonia (NH3) and Hydrogen ion (H+) producing ammonium (NH4 +) radical. Explain. CO – CONALENT BOND: The Coordinate covalent bond is formed only when an atom with an unshared pair of electrons in its valence shell donates a pair of electrons to anther atom or ion that needs a pair of electrons to acquire a stable electronic configuration. Co – ordinate Covalent Bond H | H+ H | H − N − H | H Formula H − N + NH4 + | H Q - 16) Define term metal and describe metallic bond? Metal: It is a substance, consisting of positively charged ions, fixed in a crystal lattice with negatively charged electrons moving freely through the crystal. Metallic bond: It is the combination of electrostatic attraction between the electrons and the positive nuclei of atoms. Q - 17) Explain the following properties of metals: (i) Luster (ii) Conductivity (iii) Malleability (iv) Ductility Luster (Shine): When the portable electrons in metals readily absorb light, falling upon them and move to higher energy levels. When they fall back to their original position, they emit radiation and it caused metal shine. Conductivity: As the electrons in metals are, free to move from one atom to the next they are generally good conductors of electricity and heat. Malleability: It means that metal can be easily bent or hammered into sheets. Ductility: It means that metal can be easily converted into thin wires. Q - 18) Why some metals such as sodium or potassium is soft while other are hard? Explain: Ans: It is because in Na or K valence electrons are not confined to any particular atom, instead they are free to move through crystal, and so resulting bond is relatively soft.
  • 29. 29 Q - 18) Explain the origin of dipole – dipole forces between the molecules, and give example: Dipole – dipole forces: These forces act between polar molecules that possess dipole moments. “A Dipole – dipole force, is an attractive inter-molecular force resulting from the resulting from the inter acting of the positive end of one molecule with the negative end of other.” Example: The molecule of HCl, Cl has greater electronegativity then H, a partial negative charge on chlorine (Cl) atom and partial positive charge on Hydrogen (H) atom. The (H+δ – Cl−δ) has permanent dipole moment. Q - 19) What do you mean by dispersion forces why are they also called London forces? Dispersion Forces (London Forces): These forces are the weak attractive forces between temporarily polarized atoms or molecules caused by the varying positions of the electrons during their motion about the nuclei. These forces are also called “London Forces”, because Fritz London first identified these forces in 1930. Q - 20) What is Hydrogen bonding? What type of forces, Either Intra – molecular or Inter molecular forces are present in Hydrogen bonding? Hydrogen Bonding: A hydrogen bond is a dipole – dipole attractive force that exists between two polar molecules, containing a hydrogen atom covalently bonded to an atom of F, O or N. Water is best example of Hydrogen bonding. It is denoted by (⋯). Inter – molecular forces are present in Hydrogen bonding not Intra – molecular forces. Q - 21) What do you mean by Intra – molecular and Inter molecular forces (Vander Waals’s Forces)? Intra – molecular forces: These forces hold atoms together in a molecule. For example, Water (H20) molecule consists of two hydrogen atoms, one oxygen atom join in a specific way, i.e. covalent bonds. Inter – molecular forces: They are the attractive forces between the neutral molecules, which hold them together at certain temperature. These are also called Vender Waals’s Forces.
  • 30. 30 Q - 1) Define the following term: (i) Solid (ii) Fusion (iii) Evaporation (iv) Sublimation (v) Boiling point (vi) Melting point (vii) Diffusion (viii) Brownian Movement (i) Solid: The state of matter that has definite shape as well as definite volume is known as solid. (ii) Fusion: It is the changing of state from solid to liquid by rising in temperature when heated. (iii) Evaporation: The escaping of molecules from the surface of liquid when heated is called evaporation. Or It is the changing of state from liquid to gaseous at a certain temperature. (iv) Sublimation: When solids are directly converted into gaseous state during heating this process is called sublimation. (v) Boiling point: The temperature at which the Vapor pressure of a liquid becomes equal to external pressure applied on the liquid is called boiling point. (vi) Melting point: The temperature at which the solid starts melting is called melting point. (vii) Diffusion: The spreading of substance through medium like air or liquid is called Diffusion. It is depends upon molar mass or density of substances. (viii) Brownian movement: A continuous, rapid or zigzag motion of suspended particles through the medium is called Brownian motion (movement). Q - 2) Define Kinetic molecular theory and its states of matter: According to kinetic molecular theory matter is composed of particles which are called molecules. Molecules are in motion and they possess kinetic energy. (a) Solid State: In solid state molecules are tightly packed with one another and they perform only translational motion. Due to this, molecules in solid neither slip nor slide over one other hence shape and volume of solid is definite. (b) Liquid State: In liquid state molecules are not tightly packed with one another. Their positions are not fixed and they can move in all direction, hence liquid does not have any definite shape and fixed volume. (c) Gaseous State: In gaseous state, molecules are widely separated from one another and they move freely in all directions; hence, the shape and volume of gas are not fixed. chemistry STATES OF MATTER Chapter IX SIX
  • 31. 31 Q - 3) How is solid converted into liquid? Explain it: Conversion of solid into liquid: When a solid is heated, then the kinetic energy of particles increases and becomes hot. If heating is continued then at certain temperature, the added energy becomes enough to overcome the attractive forces holding particles of solid in the fixed positions and it starts melting. At this point solid particles lose their fixed positions as well as their arrangement and thus solid is converted into liquid. Q - 4) What do mean by liquid? Describe interchange of liquid to gas: Liquid: The thing that takes the shape of vessel in which it is kept, but does not occupy total volume available is known as liquid. Interchange of liquid to gas: when liquid heats the kinetic energy of liquid molecules increases, its molecules start escaping from the surface of liquid, and Vapor pressure of liquid becomes equal to external pressure. At this point bubbles of Vapor are able to form within the interior of liquid and then rise to the surface, where they burst and release Vapor thus the liquid converts to gas. Q - 5) What is diffusion? Explain on the basis of kinetic molecular theory: Diffusion: “The spreading of a substance through medium like air or liquid is called diffusion.” In Liquid by Kinetic Molecular Theory: “According to kinetic theory freedom of liquid molecules permits diffusion to take place but the closeness of molecules and large number of collisions cause diffusion to be slow.” In Gas by Kinetic Molecular Theory: “According to kinetic theory molecules of gas are in constant random motion due to which the molecules of gas effuse throughout the vessel.” Q - 6) What id Brownian movement? Explain with example: “A continuous, rapid, zigzag motion of suspended particles through the medium is called Brownian movement” Example: Mix some Sulphur in water and stir it, after stirring filter the suspended Sulphur some of the Sulphur particles are very small and they can pass through the pores of filter paper into filtrate. Now put a drop of this filtrate on a slide and observe it under high-powered microscope. It is observed that Sulphur particles perform zigzag motion.” Explanation: According to kinetic molecular theory, water molecules are always in motion and they collide with Sulphur particles and push them in some directions where they collide with other water molecules and pushed in some other direction and this process continues.
  • 32. 32 Q - 1) Define the following term: (i) Solution (ii) Solute (iii) Solvent (iv) Suspension (v) Solubility (vi) Crystallization (vii) Molarity (M) (viii) Molality (m) (ix) Mole fraction (x) Dilute solution (xi) Concentrated solution (xii) Concentration Solution: The homogenous mixture of two or more substances with uniform composition is called solution. A solution has two components, solute and solvent. Solute: The component of solution present in smaller amount is called solute. Solvent: The component of solution present in greater amount is called solvent. When water is solvent, the solution is called aqueous solution. Suspension: A homogenous mixture that consists of visible particles, each of which contains many thousands of molecules surrounded by molecules of liquid is known as suspension. Solubility: It is define as the amount of solute in gram dissolved at a given temperature in 100 gm of the solvent. Crystallization: The process in which dissolved solute comes out of solution and forms crystals is called crystallization. Molarity: It is define as the number of moles of the solute dissolved per dm3 or 1 liter of the solution. It is denoted by M. Molality: It is define as the number of moles of the solute dissolved per kg or 1000 gm of the solvent. It is denoted by m. Mole fraction: It is a unit of concentration, defined to be equal to the number of moles of a component divided by the total number of moles of a solution. Dilute Solution: The solution that contains less amount of solute as compared to the amount of solvent is known as dilute solution. Concentrated Solution: The solution that contain greater amount of solute as compared to the amount of solvent is known as concentrated solution. Concentration: In a solution the amount of solute dissolved in a given quantity of solvent is known as its concentration. Colloid: A type of homogeneous mixture in which the dispersed particles do not settle out. Examples: butter, milk, smoke, fog, ink, paint Specific gravity: It is the ratio of the density of a substance to the density of water. Example: The specific gravity of pure water at 4 °C is 1. Specific gravity is a unit value. CHEMISTRY SOLUTION AND SUSPENSION CHAPTER IX SEVEN
  • 33. 33 Q - 2) Define Solubility and discuss the factors affecting Solubility: It is define as the amount of solute in gram dissolved at a given temperature in 100 gram of the solvent. The following are the factors affecting solubility. i.e. (i) Temperature (ii) Pressure (iii) Nature (i) : If the temperature increases the solubility of solid increases and the gases decreases in liquid. Example:  The solubility of sugar in water at 0 o C is 17.9g/100ml, whereas at 100 o C it is 48.7g/100ml.  When temperature increases the glass of cold water is warmed, bubble of air are seen inside of the glass. (ii) : The solubility of solids and liquids are not affected by pressure. But the solubility of a gas in a liquid is directly proportional to the pressure of gas, and this is called Henry’s Law. Mathematically: 𝑚 ∝ 𝑝 or 𝑚 = 𝐾𝑝 (Where “m” is the amount of gas dissolved.) Example: This effect is used in manufacture of bottled soft – drinks as coca – cola, 7 – up etc. These are bottled under a CO2 pressure slightly greater than 1 atm. When bottle are opened, pressure decreases, so the solubility of CO2 also decreases, hence the bubbles of CO2 comes out of solution. (iii) : The solubility of solute and solvent may be polar (H2O, Alcohol) and non – polar (Benzene, Carbon tetrachloride).  Polar or Ionic solutes easily dissolve in polar solvents  Non – polar solutes easily dissolve in non – polar solvents. Example: The common salt (NaCl) being ionic compound easily dissolve in polar solvent like water (H2O). But it is insoluble in non – polar solvent like Benzene or Petrol. Q - 3) Define insoluble impurity and soluble impurity: The impurity, which remains insoluble in the solvent, is called insoluble impurity. The impurity, which remains in soluble at room temperature, is called soluble impurity.
  • 34. 34 Q - 6) Define Saturated, Unsaturated and Super sturated solution: Saturated Solution: The solution which contains maximum amount of solute in a gien solvent at a specific temperature and no more solute dissolve in it is called saturated solution. In this solution there is dynamic equlibiruim between dissolved and undessolved solute. Unsaturated Solution: A solution in which the amount of the solute is less than it has the capacity to dissolve in large quantity of solvent is called unsaturated solution. Super Saturated Solution: The solution which contains greater amount of dissolved solute than that are persent in a saturated solution. It is obtained by dissolving solute in saturated solution on heating. More solute would disssolve on heating saturated solution. Q - 7) Write differences between Solution and Suspenssion: Difference Between Solution and Suspenssion Solutions:  The size of particles is between 0.1 to 1 nm.  Particles cannot be seen with low power microscope.  It is homogeneous and transparent.  Particles donot settle down.  Components cannot be seprated by filtration. Suspenssion:  The size of particles is largeer than 1000 nm.  Particles can be seen with low power microscope.  It is hectrogeneous and opaque.  Particles settle down.  Components can be seprated by filtration.
  • 35. 35 Q - 8) Name the solute and solvent in the following solution: a) Syrup b) Haze (dust in air) c) Butter (water in fat) d)Fog e) Gelies (water in fruit pulp) f) Smoke g) Sodium h)Cheese (water in fat) i) Foam (water in air) j) Mist Answer: Solution Solute Solvents a) Syrup Sucrose Water b) Haze (dust in air) Dust Air c) Butter (water in fat) Fats Water d) Fog Water vapour Air e) Gelies (water in fruit pulp) Air Water f) Smoke Fruits pulp Water g) Sodium Carbon particles Air h) Cheese (water in fat) Mercury Sodium i) Foam (water in air) Fats Water j) Mist Water vapour Air Q - 9) Explain why: i. Common salt dissolves in water but not in petrol. Ans: Common salt dissolves in water but not in petrol because like dissolve like. Common salt being polar compound easily dissolves in polar solvent like water but it is insoluble in non – polar solvent like petrol. ii. Cold drinks are bottled under a CO2 pressure greater than 1 atm. Ans: Cold drinks are bottled under a CO2 pressure greater than 1 atm because when these bottled are opened, pressure decreases, so solubility of CO2 also decreases, hence bubbles of CO2 come out of solution.
  • 36. 36 iii. 100 ml solution of KNO3 can not hold more than 37gm of KNO3 in dissolved state. Ans: 100 ml solution of KNO3 can not hold more than 37gm of KNO3 in dissolved state because this amount of KNO3 is required in making saturated solution. Q - 10) What do you mean by the percentage concentration and write some example of suspenssion in daily life. Percentage Concentration: The percentage concentration is based on mass (M) and volume (V) of the components solute and solvent in the solution. It has four ways to express: (i) Percentage in 𝑚 𝑚 % ( 𝑚𝑎𝑠𝑠 𝑚𝑎𝑠𝑠 ) % (ii) Percentage in 𝑚 𝑣 % ( 𝑚𝑎𝑠𝑠 𝑣𝑜𝑙𝑢𝑚𝑒 ) % (iii) Percentage in 𝑣 𝑚 % ( 𝑣𝑜𝑙𝑢𝑚𝑒 𝑚𝑎𝑠𝑠 ) % (iv) Percentage in 𝑣 𝑣 % ( 𝑣𝑜𝑙𝑢𝑚𝑒 𝑣𝑜𝑢𝑚𝑒 ) % Examples (i) 05% ( 𝑚 𝑚 ) Solution means solute 5g in 95gm solvent. (ii) 10% ( 𝑚 𝑣 ) Solution means solute 10g in solution 100cm3 (iii) 05% ( 𝑣 𝑚 ) Solution means solute 5 cm3 in solution 100g. (iv) 15% ( 𝑣 𝑣 ) Solution means solute 15cm3 in solution 85cm3 solvent. The examples of suspennsion in daily life: i. Smoke A suspension of the particles of carbon in gas or air. ii. Mud (smile) A Suspension of the particles of solid in small quantity of liquid. iii. Foam (forth) A suspension of fine particles of a gas in a liquid. iv. Emulsion A suspension of droplets of one liquid into another in which it is not soluble.
  • 37. 37 Q - 1) Define the following term: (i) Electrochemistry (ii) Electrolytes (iii) Non – Electrolytes (iv) Electrolysis (v) Electroplating (vi) Electrochemical equivalent (vii) Ampere (viii) Coulomb (ix) Faraday (x) Batteries Electrochemistry: The branch of chemistry that deals with the relationship between electricity and chemical reactions i.e. inter conversion of electric energy and chemical energy is defines as Electrochemistry. Electrolyte: The chemical compound, which conducts electricity in molten condition or by its aqueous solution with chemical change, is called an Electrolyte. Examples: Hydrochloric acid (HCl), Sodium hydroxide (NaOH), Sodium Chloride (NaCl) etc. Non - electrolytes: The compounds which do not conduct electricity in molten or in aqueous solutions are called Non – electrolytes. Ex: Sugar, petrol, benzene etc. Electrolysis: A process in which movements of the ions take place towards their respective electrodes to undergo chemical changes under the influence of an applied electric field is known as Electrolysis. Electroplating: The process of electrolysis that is used to coat one metal onto another metal is known as Electroplating. It protects baser metal from corrosion and makes them more attractive. Electrochemical equivalent: The weight of substance deposited or liberated when one coulomb of electric charge is passed through an electrolyte is known as Electrochemical Equivalent. It is denoted by Z, and its SI unit is kg/coulomb. Ampere: It is the basic unit of current that passes through a circuit for one second that can liberate 0.001118g of Ag from silver nitrate (AgNO3) solution. Coulomb: It is the basic unit of charge and is define as the quantity of charge when one ampere of current is passed for one second. C = Ampere (A) × time (s) Fraday: It is defining as the quantity of charge, which deposits or liberates exactly one-gram equivalent of a substance. 1F = 96500 coulombs. Batteries: The devices to produce electricity by the chemical reactions are known as batteries. A battery is an assembly of two or more voltaic cells. Electron affinity reflects the ability of an atom to accept an electron. The energy change occurs when an electron is added to a gaseous atom. Atoms with stronger effective nuclear charge have greater electron affinity. Ex: H(g) + e - → H - (g); ΔH = -73 kJ/mol. chemistry ELECTROCHEMISTRY Chapter IX EIGHT
  • 38. 38 Q - 2) State and explain Faradays first law of electrolysis: : It states that the amount of any substance deposited of liberated at an electrode during electrolysis is directly proportional to the quantity of current passed through the electrolyte. : If “w” is the weight or amount of a substance deposited or liberated and “A” ampere of current is passed of “t” seconds, then according to the law: w ∝ A × t w = Z ×A ×t Where “𝑍” is a constant; known as electrochemical equivalent. Q - 3) State and explain Faradays Second law of electrolysis: : It states that the masses of different substances deposited or liberated when same quantity of current is passed through different electrolytes, connected in series are proportional to their chemical equivalent masses. : If we consider three different electrolytes, AgNO3, CuSO4 and Al(NO3)3 solutions, connected in series and same amount of current is passed through them. Then the masses of Ag, Cu and Al deposited on their respective electrodes would be directly proportional to their equivalent masses. If96500coulombsofelectricchargeispassedthenthemassesofAg,CuandAlwould be 108g, 31.75g and 9g respectively. Formula: 𝐸𝑞𝑢𝑖𝑣𝑎𝑙𝑒𝑛𝑡 𝑚𝑎𝑠𝑠 𝑜𝑓 𝑎𝑛 𝑒𝑙𝑒𝑚𝑒𝑛𝑡 = 𝐴𝑡𝑜𝑚𝑖𝑐 𝑚𝑎𝑠𝑠 𝑜𝑓 𝑒𝑙𝑒𝑚𝑒𝑛𝑡 𝑉𝑎𝑙𝑒𝑛𝑐𝑦 𝑜𝑓 𝑡ℎ𝑒 𝑒𝑙𝑒𝑚𝑒𝑛𝑡 Q - 4) Write down the relationship between equivalent mass and electrochemical equivalent: If “e” is the gram equivalent mass and “Z” is Electrochemical Equivalent (E.C.E) then we can write it as: 𝑒 = 𝐹 × 𝑍
  • 39. 39 Q - 5) What do mean by Primary cell also describes the construction and working of Dry cell with diagram: Primary Cell: The cell, which is used to convert chemical energy into electrical energy and that, has irreversible characteristics, is known as Primary Cell or Dry Cell. : i. It has out zinc (Zn) vessel, which acts as anode, and inert carbon (graphite) rod, which acts cathode. ii. A mixture of manganese dioxide (MnO2) and carbon powder surrounds the graphite rod. iii.The electrolyte is a moist paste of ammonium chloride (NH4Cl) and zinc chloride (ZnCl2). iv.The upper top position is sealed with wax and a copper cap is fitted on the top carbon rod to make the electrical contact. v. The whole cell is covered with a safety cover. : When a metallic wire Zn connects zinc and graphite electrodes is oxidized to form Zn2+ ions, which pass into the wet paste leaving behind electrons on the Zn container the electrons, move from Zn electrode to carbon electrode through the external circuit. Q - 6) Describes the construction and working of Daniel cell with diagram: Daniel Cell: It is the simplest of the Galvanic or Voltaic cell which is used to convert chemical energy into electrical energy. : i. It consists of two half-cells, i.e. Zinc rod and Copper rod. ii. Zinc rod is dipped in 1M ZnSO4 solution and Copper rod is dipped in 1M CuSO4 solution. iii. The two half-cells electrodes are connected together to form a complete cell. iv. A porous partition or a salt bridge separates the two half-cells. v. The electrodes are connected externally through a volt Meter by metal wire.
  • 40. 40 : When the connection are fully made. The cell starts producing electric current at once. (i) Zn undergoes oxidation to form Zn2+ ions by the loss of 2e– two electrons to go into ZnSO4. And acts as anode or negative electrode. (ii) The electrons, which are free at Zn electrode, travel through the wire externally to Cu electrode moreover accepted by CuSO4 solution. (iii) Cu2+ ions undergo reduction to deposit copper metal at cu electrode, which acts as cathode, or positive electrode. (iv) In this process Zn – electrode dissolves in the solution of ZnSO4 and reduces in size, while Cu – electrode grows in size due to the deposition of Cu – metal. : 1.10 volt. At Zn Electrode (Anode) Zn(s)  Zn2+ (aq) + 2e– At Cu Electrode (Cathode) Cu(aq) + 2e– Cu(s) Zn(s) + Cu2+ (aq)  Zn2+ (aq) + Cu(s) The Diagram Of Daniel Cell VOLT METERZinc Anode SO4 2- Cu4 2- SO4 2- Copper Cathode Salt Bridge ZnSO4 Solution CuSO4 Solution Zn4 2-
  • 41. 41 Q - 7) What happens when electric current is passed through acidulated water? Give the reaction at two electrodes and mention the products at cathode and anode. Ans: When electric current is passed through acidulated water it conducts electricity. H2O(l) H3O+ (aq) + OH– (aq) 1. The (Hydronium) positive ions (H3O+ ) move towards cathode, gain electrons at cathode and get neutralized to liberate H2 gas. 2. The (Hydroxide) negative ions (OH– ) move towards anode, lose electrons at anode and get neutralized to liberate O2 gas. The reaction at cathode: H3O+ (aq) + 2e– H2 (g) + 2H2O(l) The reaction at anode: 4OH – (aq) O2 (g) + 2H2O(l) + 4e– The overall reaction: 2H2O(l) 2H2 (g) + O2 (g) Result: It is observed that on electrolysis of water we get two volumes of hydrogen gas for each volume of oxygen gas. Humphrey Davy who first did the electrolysis of water and confirmed the formula of water is H2O. Q - 8) Predict the net electrolysis reaction when molten NaCl is electrolyzed. Or describe the electrolysis of molten Sodium Chloride. Ans: Sodium Chloride, NaCl (Salt) does not conduct electricity in the solid state because its ions are held together tightly in a regular lattice arrangements and cannot move but when it is electrolyzed or melted the ions are freed from their lattice state and can move freely to conduct electricity. Procedure of Making Electrolyte: i. Some fused NaCl is taken in an electrolytic cell i.e. in a glass vessel. ii. Two platinum rods (electrodes) are dipped into the fused salt (NaCl) and connected to the battery outside the cell by wires. iii. The battery has two terminal negative terminal as cathode and positive terminal as anode. Acid
  • 42. 42 Working of the cell: When an electrical potential is passed through the Molten Sodium Chloride (NaCl) salt, electrolysis starts. i. The positive ions i.e. captions (Na+) are attracted towards cathode. ii. The negative ions i.e. anions (Cl–) are attracted towards anode. iii. At both electrodes, chemical reactions take place. The reaction at cathode: Na+ (l) + e– Na(l) The reaction at anode: 2Cl – (g) Cl2 (g) + 2e– The overall reaction: 2Na+ (l) + 2Cl– (l) 2Na(l) + Cl2 (g) Result:ItisobservedthatonelectrolysisofmoltenSodiumChloride(NaCl) wegetsodium metal (Na) at cathode and Cl2 gas is liberated at anode. Q - 9) Difference between Primary cell and Secondary cell. Primary Cell: i. It is used to convert chemical energy into electrical energy and it has dense material. ii. It becomes dead after sometimes and cannot be reused. (But nowadays possible) iii. Example Dry Cell and mercury cell. Secondary Cell: i. It is also used to convert chemical energy into electrical energy that has liquid material. ii. It can be reused over and again by recharging. iii. Example Lead Storage battery and nickel – cadmium storage cell. Battery Cl – Na+ Anode (+) Cathode (–)
  • 43. 43 Q - 10) Describe the process of nickel plating. Nickel Plating: The process of dropping a thin layer of nickel metal over anther metal object with the help of electrolysisiscallednickelplating. Procedure: i. A cell for electroplating of nickel consists of a piece of pure nickel as anode and the spoon or any other object as cathode. ii. A solution of nickel Sulphate (NiSO4) is used as the electrolyte in the electrolytic cell. iii. In addition, both electrodes (anode and cathode) join externally with a battery. iv. On passing electric current, the anode (Ni) dissolves in the solution and forms Ni2+ ions. v. Then Ni2+ ions from the solution move towards cathode. They gain electrons and reduce to Ni metal on the surface of spoon (cathode). The reaction at cathode: Ni2+ (aq) + 2e– Ni(s) The reaction at anode: Ni(s) Ni2+ (aq) + 2e– The overall reaction: Ni(s) Ni(s) Result: It is observed that the net reaction is simply the transfer of Ni as Ni2+ through NiSO4 solution towards the cathode i.e. spoon and get it coated with Ni metal on the surface. Q - 11) Describe the process of Chromium plating. Chromium Plating: Theprocessofdroppingathinlayerofchromiummetaloveranthermetalobjectwiththehelpof electrolysisiscalledchromiumplating. Procedure: i. A cell for electroplating of chromium consists of a piece of pure chromium as anode and the spoon or any other object as cathode. ii. A solution of chromium Sulphate Cr2 (SO4)3 is used as the electrolyte in the electrolytic cell. iii. In addition, both electrodes (anode and cathode) join externally with a battery. iv. On passing electric current, the anode (Cr) dissolves in the solution and forms Cr3+ ions. v. Then Cr3+ ions from the solution move towards cathode. They gain electrons and reduce to Cr metal on the surface of spoon (cathode). The reaction at cathode: Cr3+ (aq) + 3e– Cr(s) The reaction at anode: Cr(s) Cr3+ (aq) + 3e– The overall reaction: Cr(s) Cr(s) Result: Itis observed that the net reaction is simply the transfer of Cr as Cr3+ through Cr2 (SO4)3 solution towards cathode i.e. spoon and get it coated with Cr metal on the surface.
  • 44. 44 Q - 12) What is the function of salt bridge or porous partition in an electrochemical cell? Function of salt bridge: The function of salt bridge or porous partition in electrochemical cell is to prevent the mixing of two solutions and it allows ions to move through one part to another. Q - 13) Predict what would be formed (i) at the anode and (ii) at the cathode when each of the following molten salts is electrolyzed using inert electrodes. a) NaCl b) MgBr2 c) CaCl2 Salts At cathode At anode a) NaCl Na (l) Cl2 b) MgBr2 Mg (l) Br2 c) CaCl2 Ca (l) Cl2 Q - 14) Write down the uses of electrolysis. Uses of electrolysis: I. It is used for the extraction of certain metals from their ores. II. It is also used in electroplating of metal. III. It is also used to purify many metals Q - 15) Define electrochemical cell. Electrochemical cell: The cell, which is used to convert chemical energy into electrical energy, is called electrochemical cell and it is generally known as Galvanic or Voltaic cell. Examples: The simplest of the Galvanic or Voltaic cells is Daniel cell. Q - 16) Which of the following pairs of terms have the same meanings and which have different meanings? (i) Voltaic cell and Galvanic cell. (ii) Cell and battery (iii) Electrolytic cell and electrochemical cell Ans: The pair of Voltaic and Galvanic cell has the same meanings because two pioneers in electricity Luigi Galvani and Alexandra Volta exposed it. The other two pairs are different in meanings because they have different properties and processing procedures.
  • 45. 45 Q - 1) Define the following term: 1) Salt 2) pH 3) Hydrolysis 4) Titration 5) Neutralization 6) Basicity of Acid 7) Double salts 8) Standard solution 9) Acidity of base Salt: A salt is ionic compound produced when a base neutralizes an acid. Example NaCl, KNO3. pH: It is define as the negative logarithm of the hydrogen ion H+ or H3O+ ions concentration in moles per liter. Hydrolysis: The complete dissociation of salt into acid or base in aqueous solution is known as hydrolysis. Titration: It is the chemical process by which, we can determine the concentration of unknown solution. Neutralization: A reaction in which an acid and a base form an ionic compound (salt and water) is called neutralization. Basicity of acid: The number of replaceable or ionize hydrogen (H+) ions present in molecule of an acid is called basicity of acid. Double salts: The crystalline compounds, which are obtained when two specific salts are crystallized together, are known as double salts. Standard solution: The solution whose concentration is known is called standard solution. Acidity of base: The number of replaceable or ionize hydroxide ions (OH–) present in molecule of base is called acidity of base. Q - 2) Write down the importance of pH. Importance of pH: I. pH plays great role in the field of Biology. II. It is also important for water treatment and food processing. III. It also plays an important role in corrosion control and soil conditioning. IV. It has some important in the process of electroplating. The pH values of several biological fluids. Lemon juice 2.3 Human blood 7.35 – 7.45 Vinegar 2.8 Cow’s milk 6.5 Tomato juice 4.2 Human urine 5.0 – 7.0 Salvia 7.0 Egg white 7.8 chemistry ACIDS BASES AND SALTS Chapter IX NINTH
  • 46. 46 Q - 3) State Bronsted Lowry theory of acids and bases and explain it with example of HCl and NH3. : “He states that an acid is a substance having a tendency to donate one or more protons while base is substance having a tendency to accept (add) protons.” Bronsted – Lowry acid = A substance that can donate (H+ ). Bronsted – Lowry base = A substance that can accept (H+ ). : Lowry takes hydrochloric acid (HCl) and water (H2O) is protons donors and act as Bronsted-Lowry acids whereas (H2O) and ammonia (NH3) are protons acceptors and are known as Bronsted Lowry bases. ACID ACIDBASE BASE H – ×O+H×O×H HN H × × H × HN H × H × H × + H × C l + H×0 H × × C l – H0 H × H × + PROTON DONOR +
  • 47. 47 Q - 4) State Lewis concept of acids and bases and explain it with example: : He states that an acid is any species (molecules or ions) which can accept a pair of electrons, and base is any species that can donate a pair of electrons. Lewis acid = An electron pair acceptor. Lewis base = An electron pair donor. : When ammonia (NH3) reacts with proton (H+ ) to form ammonium ion (NH4 + ) in which the Nitrogen of NH3 donates a pair of electrons whereas the (H+ ) accepts that pair of electrons for bond formation. Q - 5) What is Arrhenius theory of acid and bases why is the Arrhenius theory not satisfactory for acids and bases: : He states that an acid is any substance that yields hydrogen ions (H+ ) in water and a base is a substance that yields hydroxide ions (OH– ) in water. Arrhenius acid = HCl and H2SO4 HCl(aq) H+ (aq) + Cl– (aq) Arrhenius base = NaOH and KOH NaOH(aq) Na+ (aq) + OH– (aq) His theory does not satisfy for acids and bases because they apply only to water solution and it does not account for the basicity of ammonia (NH3) that doesn’t contain (OH) group.
  • 48. 48 Q - 6) Write down the physical and chemical properties of Acids and Bases: Acid is a substance that gives H+ or H3O+ in aqueous solution. Base is a substance that gives OH – or in aqueous solution. : : 1. They have a sour taste. 2. Change the color of litmus form blue to red. 3. Strong acids can harm human, things etc. 4. They conduct electricity in aqueous solution. 1. They have bitter taste. 2. Changethecoloroflitmusfromredtoblue. 3. They have slippery touch. 4. They also conduct electricity in aqueous solution. : 1. When acids react with metals like (Zn, Mg and Fe) they produce H2 gas and salts. Zn(s) + 2HCl(aq) ZnCl2(s) + H2(g) Mg(s) + 2HCl(aq) MgCl2(s) + H2(g) 2. When acids react with carbonate and bicarbonates they produce CO2 gas. CaCO3(s) + 2HCl(aq) CaCl2(aq)+ H2O(l) + CO2(g) NaHCO3(s)+ HCl(aq) NaCl(aq) + H2O(l) + CO2(g) 3. When acids react with oxides and hydroxides of metals they produce salts and water. FeO(s) + 2HCl(aq) FeCl2(aq) + H2O(l) Fe(OH)3(s)+ 3HCl(aq) FeCl3(aq) + 3H2O(l) 4. When acids react with bases they produce salts and water. NaOH(aq) + 2HCl(aq) NaCl(s) + H2O(l) Ca(OH)2 + 2HCl(aq) CaCl2(s) + H2O(l) : 1. Bases react with fats to form soap. 2. When bases react with metals or nonmetals they dissolve them and produce H2 gas. 2Al + 2NaOH(aq) + 2H2O(l) 2NaAlO2 + 3H2 3. When bases react with acid they produce water and salts. NaOH(aq) + 2HCl(aq) NaCl(s) + H2O(l) Ca(OH)2 + 2HCl(aq) CaCl2(s) + H2O(l) 4. When bases react with heavy metal salts they form their salt solutions. 3NaOH + FeCl3 Fe(OH)3 + 3NaCl 2NaOH + MgCl2 Mg(OH)2 + 2NaCl
  • 49. 49 Q - 7) What is salt and explain its four examples: A salt is ionic compound produced when an acid is neutralized by a base. For example sodium hydroxide (NaOH) neutralizes hydrochloric acid (HCl) to form sodium chloride (NaCl) and water (H2O). NaOH(aq) + 2HCl(aq) NaCl(s) + H2O(l) There are four kinds of salts as following: (i) Normal Salt (ii) Acidic Salt (iii) Basic Salts (iv) Double Salt (i) Normal Salt: Salts which are formed by the complete neutralization of an acid by a base e.g. (NaCl, NaNO3, K2SO4) etc. are known as normal salts and these do not have replaceable hydrogen ions (H+ ) or hydroxyl groups (OH– ). (ii) Acidic Salt: Salts, which are formed by the partial neutralization of an acid by a base e.g. (KHCO3,NaHSO4)etc.areknown asacidicsalts andthese, havereplaceablehydrogen atoms(H+ ). NaOH(aq) + H2SO4(aq) NaHSO4(s) + H2O(l) (iii) Basic Salt: Salts which are formed by the partial neutralization of an base by an acid e.g. Mg(OH)Cl, Zn(OH)Cl etc. are known as basic salts and these have replaceable (OH– ) hydroxyl groups. Mg(OH)2(aq) + HCl(aq) Mg(OH)Cl(s) + H2O(l) (iv) Double Salt: The crystalline compound which are obtained, when two specific salts are crystallized together are called double salts. These have definite chemical composition. Such as (i) Potash Alum K2SO4 Al2(SO4)3 24H2O (ii) Mohr’s Salt FeSO4 (NH4)2 SO4 (iii) Chrome Alum K2SO4 Cr2(SO4)3 24H2O (iv) Carnalities KCl MgCl26H2O Q - 8) What is chemical reaction define and give one example each of the following (i) Neutralization (ii) Hydrolysis Chemical Reaction: The process in which chemical change in the nature and composition of some substance occurs is called a chemical reaction. Neutralization: The process in which two reactants acid and base reacts with each other to form salt and water is called neutralization. NaOH(aq) + 2HCl(aq) NaCl(s) + H2O(l) Hydrolysis: The process in which acid or base reacts with water to form either weak acid and strong base or strong acid and weak base is called hydrolysis. FeCl3 + 3H2O Fe(OH)3 (Weak Base) + 3HCl (Strong Acid) Na2CO3 + 2H2O 2NaOH (strong Base) + H2CO3 (Weak Acid)
  • 50. 50 Q - 9) Explain with illustration that what are strong acid and base and weak acid and base. Also write formulas of them: Strong Acid: An acid that produces large number of H+ ions in aqueous solution is said to be a strong acid and it is almost completely dissociated. Weak Acid: An acid that produces small number of H+ ions in aqueous solution is said to be a strong acid it is only partially dissociated. Strong Base: A base that produces large number of OH– ions in aqueous solution is said to be a strong acid and it is almost completely dissociated. Weak Base: A base that produces small number of OH– ions in aqueous solution is said to be a strong acid it is only partially dissociated. Formulas of acid and bases Strong acid Strong base Hydrochloric Acid HCl Sodium Hydroxide NaOH Nitric Acid HNO3 Potassium Hydroxide KOH Sulphuric Acid H2SO4 Barium Hydroxide Ba(OH)2 Weak acid Nitrous Acid HNO2 Carbonic Acid H2CO3 Phosphoric Acid H3PO4 Acetic Acid CH3 – COOH Hydro fluoric Acid HF Formic Acid HCOOH Weak base Ammonium Hydroxide NH4OH Magnesium Hydroxide Mg(OH)2 Beryllium Hydroxide Ba(OH)2 Q - 10) Sulphuric acid (H2SO4) is strong acid. H2SO4 - is weak acid account for the difference in strength: H2SO4: H2SO4 is strong acid because it ionizes completely in aqueous solution and produces two H+ ions. While HSO4 – : HSO4 – is weak acid because it ionizes partially in aqueous solution and produces only one H+ ion. Q - 11) Give an example of mono-parotic, di-parotic and tri-parotic acid: Mono-practice acid = HCl Di-parotic acid = H2SO4 Tri-parotic acid = H3PO3
  • 51. 51 Q - 12) Define acidic, basic and neutral solutions in terms of H+ ion concentration. Indicate whether each of the following solution will be acidic, basic or neutral: a) Strong acid and strong base Ans: Neutral solution b) Strong acid and weak base Ans: Acidic solution c) Weak acid and strong base Ans: Basic solution Acidic solutions: It is the solution in which number of hydrogen ions (H+ ) increase is said to be acidic solution. Basic solutions: It is the solution in which number of Hydroxide ions (OH– ) increase is said to be basic solution. Neutral solutions: The solution in which number of hydrogen ions (H+ ) and hydroxide ions (OH– ) become equal is said to be Neutral solution. Q - 13) Define molarity. What is molar solution? Molarity: It is defined as the number of moles of solute dissolved per 1 litter or 1dm3 solutions of a solution and denoted by M. Molar solutions: The solution which contains 1 mole of solute in 1dm3 (1liter) of solution is known as molar solution. Q - 14) Define term Amphoteric and give example? Amphoteric: A substance (such as water) that can behave as both acid and base is said to be an amphoteric substance. Example: HCl(aq) (Acid) + H2O(l) (Base) H3O+ (aq) (Acid) + Cl– (aq) (Base) H2O(l) (Acid) + NH3(aq) (Base) NH4(aq) (Acid) +OH– (aq) (Base) Q - 15) Give an equation to show the dissociation of water? Dissociation of water equation: A proton from one water molecule is transferred to another water molecule, leaving behind (OH– ) ion and forming H3O+ ion. 𝑀 = 𝑚 𝐹𝑀 × 1000 𝑣
  • 52. 52 Q - 16) Define pH? Explain why a natural solution is said to have a pH of 7? pH: 1. It is defined as the negative logarithm of hydrogen ion (H+ ) or (H3O+ ) concentration in mole per litter. 2. A natural solution is said to have a pH of 7 because in any natural solution the H+ ions concentration is1 × 10−7 M, having pH of 7. Mathematically: Titration results by pH pH = − log H+ pH = − log 10−7 pH = −(−7) pH = 7 Indicator Colour in acid pH – range Color in base Methyl orange Red 3 – 5 Yellow Litmus Red 6 – 8 Blue Phenolphthalein Colorless 8 – 10 Pink (Red) Q - 17) Describe the commercial preparation and uses of salts? 1. Commercial Preparation of SODIUM CARBONATE: (Na2CO3.10H2O) (Na2CO3) Sodium carbonate is commercially prepared by “Solvay process” or Ammonia soda process. The raw materials are limestone (CaCO3), sodium chloride (NaCl) and ammonia (NH3) and water (H2O). Step of process (i) Lime stone is heated to yield calcium oxide (CaO) and the CO2 gas. (ii) This CO2 is passed into aqueous solution of ammonia, and the ammonium bicarbonate (NH4HCO3) is produced. (iii) This (NH4HCO3) reacts with aqueous cold solution of NaCl at 15 0 C, called Brine to yield (NH4HCO3) which is not soluble at low temperature 15 0 C and this precipitates out. (iv) This (NaHCO3) on heating yields sodium carbonate (Na2CO3). Anhydrous sodium carbonate (Na2CO3) is known as soda – ash and sodium carbonate dehydrate (Na2CO3 10H2O) is commonly known as washing soda. Chemical equations H CaCO3 (s) CO(s) + CO2(g) NH3(g) + CO2 (g) + H2O(l) NH4HCO3(aq) At 15 0C brine NH4 + HCO3 – (aq) + Na+Cl– NH4 + Cl– (aq) + NaHCO3(s) H 2NaHCO3(s) Na2CO3(s) + H2O(l) + CO(g)
  • 53. 53 Uses of sodium carbonate (Na2CO3) (i) It is used in the softening of water. Sodium carbonate finishes carbonate ion to precipitate calcium and magnesium ions. (ii) It is used as cleaning agent and in making of soap. Detergents and paper. (iii) It is used in making ordinary glass (Na2SiO3), which is used in bottles. 2. Commercial preparation of Sodium hydrogen Carbonate: (NaHCO3) (Baking Soda) “Solvay Process” forms (NaHCO3) sodium hydrogen carbonate or baking soda but mostly it is prepared by passing the stream of CO2 through concentrated aqueous Na2CO3 solution. Na2CO3 (s) + CO2(g) + H2O(l) 2NaHCO3 (s) Uses of sodium hydrogen carbonate (NaHCO3) (i) It is used in the preparation of baking powder. (ii) It is also used in the preparation of effervescent drinks and fruit salts. (iii) It is also used in medicines to remove acidity of stomach (i.e. as Antacid) (iv) It is also used in fire extinguishers. 3. Commercial preparation of Copper Sulphate: (CuSO45H2O) (Blue Vitriol) (CuSO4) copper Sulphate or cupric Sulphate, which is also known as blue vitriol, or blue stone. (i) Reacting copper scraps with dilute Sulphuric acid in the presence of air prepare it. 2Cu (s) + 2H2SO4(aq) + O2 2CuSO4 (aq) + 2H2O (ii) It can be also prepared by the treatment of CuO or CuCO2 with dilute Sulphuric acid (H2SO4). CuO + H2SO4(aq) (dilute) CuSO4 + H2O CuCO3 + H2SO4(aq) (dilute) CuSO4 + H2O + CO2 Uses of sodium hydrogen carbonate (NaHCO3) (i) It is used in textile (mordant), tanning, electric batteries, hair dyes and in electroplating. (ii) It is used as germicide, insecticide, preservative for wood, paper pulp. (iii) It is used in calico printing, making synthetic rubber and copper salts e.g. steels, green paint. (iv) It is used in paint and varnish industries. (v) It is also used to kill fungus and molds.
  • 54. 54 4. Commercial preparation of Magnesium Sulphate: (MgSO47H2O) (Epsom - Salts) (MgSO4) Magnesium Sulphate is also known as “Epsom Salts”. (i) It is prepared by reacting H2SO4 with magnetite (MgCO3). MgCO3 (magnesite) + H2SO4(aq) MgSO4 + 2H2O + CO2 (ii) It is also prepared by reacting H2SO4 with dolomite (MgCO3 CaCO3). MgCO3 CaCO3 (dolomite) + 2H2SO4(aq) MgSO4 + CaSO4+ 2H2O + CO2 (iii) It is nowadays prepared by heating kieserite (MgSO4H2O) under pressure with water (H2O). MgSO4 H2O (kieserite) + 6H2O MgSO4 7H2O (Epsom salt) Uses of magnesium Sulphate (MgSO47H2O) (i) It is used as a mild purgative in medicines. (ii) It is used in dying and tanning process. (iii) It is used in fire proof fabrics. (iv) It is used in paper industry. (v) It is used in manufacture of ceramics, glazed tiles and match boxes. 5. Commercial preparation of Potash Alum: (K2SO4Al2SO424H2O) (Ordinary Alum) Potash Alum is known as ordinary salts. It is prepared by reacting K2SO4 and Al2SO4with water. When equ – molecular quantities of K2SO4 and Al2SO4 are dissolved in water and the solution is allowed to evaporate. Crystals of K2SO4 Al2SO4 24H2O which is called Ordinary Alum or Potash Alum are separated out. K2SO4 + Al2SO4 + 12H2O K2SO4 Al2SO4 24H2O Uses of potash alum (K2SO4 Al2SO4 24H2O) (i) It is used as antiseptic and as a mouth wash. (ii) It is used in sizing paper and tanning leather. (iii) It is used in purifying water. (iv) It is used in dying as mordant to fix insoluble dye to fiber. (v) It is used in medicine. Q - 18) Define acid – base titration, standard solution and equivalence point? Standard Solution: A solution whose molarity or strength is known is called standard solution.
  • 55. 55 Acid – Base Titration: In Acid – Base titration, a solution of known concentration (base) is added gradually to a solution of unknown concentration (acid) so as to determine the concentration of unknown solution. Equivalence point or End Point: The point at which reaction is completed is called the end point. At this point acids and base becomes equivalent to neutralize each other. This point is also known as the Equivalence Point. Q - 19) Identify the following as a weak or strong acids or bases? (i) NH3 (ii) H3PO4 (iii) LiOH (iv) HCOOH (Formic acid) (v) H2SO4 (vi) H2CO3 (vii) Ba(OH)2 Answer: Strong Acid: Strong Base H2SO4 Ba(OH)2 Weak Base Weak Acid LiOH, NH3 H2CO3 , H3PO4, HCOOH (Formic acid) Q - 20) Describe clearly how a solution of HCl could be titrated with a solution of NaOH? Step of Titration: i. Fill the burette with the given solution of NaOH. Read and record the initial reading then let the burette drain into beaker down to the zero mark before using it. ii. Pipette out 10cm3 of HCl in a conical flask. iii. Add one or two drops of phenolphthalein indicator. iv. Add slowly the NaOH solution, from the burette into the conical flask with constant shaking. v. Stop adding the NaOH solution when the mixture in the titration flask becomes light pink. vi. Record the final burette reading. vii. Take more readings by repeating experiment in same manner. Information required for calculation of titration: The following information is required for calculation of titration. i. The volumes of the solution of acids or base must be known. ii. The concentration of one of the solution also must be known. iii. The balanced chemical equation for the reaction.
  • 56. 56 Q - 1) Define the following term: 1. Heat content (Enthalpy) 2. Heat of Neutralization 3. Thermochemistry 4. Exothermic reaction 5. Endothermic reaction Heat content (Enthalpy): The energy given out or absorbed at constant pressure is called enthalpy. It is denoted by H. Heat of Neutralization: The amount of heat released during a neutralization reaction in which 1 mole of water is formed is called as the heat of neutralization. Thermochemistry: The branch of chemistry that deals with the study of heat changes in chemical reactions is called Thermochemistry. Exothermic reaction: It is the chemical change during which heat is given out or released. It is denoted by –H. Endothermic reaction: It is the chemical change during which, heat is taken in or absorbed. It is denoted by +H. Q - 2) Define a thermochemical reaction. What is meant by enthalpy of a reaction? Also Define Exothermic reaction with two examples: Thermochemical reaction: The chemical reactions during which material changes are accompanied with change in heat energy are called Thermochemical Reaction. Enthalpy: The heat given out or absorbed at constant pressure is called enthalpy. It is denoted by H. Exothermic reaction: The chemical change in which heat is given out or released is known as exothermic reaction. It is denoted by –H. Example of exothermic reaction: i. The combustion of coal in air is the example of exothermic reaction. 393.7 kilo Jules per mole of heat energy is released when 1 mole of coal is burnt in 1 mole of O2 to produced 1 mole of CO2. C(s) + O2 (g) CO2(g) – H = -393.7 kj/mol ii. Burning of methane in presence of oxygen is 2nd example of exothermic reaction. CH4(s) + 2O2 (g) CO2(g) + 2H2O – H = -890 kj/mol chemistry CHEMICAL ENRGETICS Chapter IX ten
  • 57. 57 Q - 3) Define Endothermic reaction and also give two examples Endothermic Reaction? Endothermic reaction: The chemical change in which heat is taken in or absorbed is known as endothermic reaction. It is denoted by +H. Example of endothermic reaction: i. The decomposition of water into Hydrogen and Oxygen is the example of endothermic reaction. 286 kilo Jules/mole of heat energy is absorbed when 1 mole of H and ½mole of O decomposed. H2O(l) H2(g) + ½ O2 (g) + H = +286 kj/mol ii. Combination of Nitrogen (N) and Oxygen (O) is 2nd example of endothermic reaction. ½N2(g) + ½O2 (g) NO (g) + H = +90.25 kj/mol Q - 4) Define Heat of Neutralization: what would be the value of reaction of neutralization when strong acid reacts with strong base? Heat of Neutralization: The amount of heat released during a neutralization reaction in which one mole of water (H2O) is formed is called as the Heat of neutralization. The value of reaction of neutralization would be almost same when strong acid react with strong base. Q - 5) How is food warm by exothermic reaction? Warming Food by exothermic reaction: The pouch that contains the food is attached to flameless radiation heater. The heater contains chemicals that react with water to produce heat. When the pouch is placed in a bad and water added temperature of the food reaches to 60 0 C in about 15 minutes. Mg(s) + 2H2O(l) Mg(OH)2(s) + H2 (g) H = –3.53 kj/mol THE END OF FIRST BOOK PART 1 Part 2 coming soon
  • 58. 58 Q - 1) Define Hydrogen and write down uses, physical and chemical properties of Hydrogen: Hydrogen: A colorless, highly flammable gaseous element, the lightest of all gases and the most abundant element in the universe, used in the production of synthetic ammonia and methanol, in petroleum refining, in the hydrogenation of organic materials, as a reducing atmosphere, in ox hydrogen torches, and in rocket fuels. Atomic number 1; atomic weight 1.00794; melting point -259.14°C; boiling point -252.8 °C; density at 0 °C 0.08987 gram per liter; valence 1. Physical Properties of Hydrogen: 1. It is a colorless, odorless and tasteless gas. 2. It is insoluble in water and highly inflammable gas and burns with blue flame. 3. Its bond dissociation energy is 104 k. cals/mol and ionization energy is 13.54 ev. 4. It liquefies at -252 °C and freezes at -259 °C. 5. Electro negativity of hydrogen is 2.1. Uses of Hydrogen: 1. It is used in manufacture of fertilizers. 2. It is used in manufacture of tungsten bold filaments and vegetable ghee. 3. It is used for purification of metals. 4. It is used in the preparation of oxy-hydrogen flame, which is used in welding due to production of high temperature. 5. It is used in the preparation of chemicals like NH2, CH3OH, etc. Chemical Properties of Hydrogen: 1. Decomposition of molecular Hydrogen (H2) Molecular hydrogen contains stable covalent bonding and is relatively inert at ordinary conditions. Its bond dissociation energy is 104 K.cals/mol. H—H H + H H = 104 K.cals/mol. i.e +435 KJ/,mol. Q. Show hydrogen is a good reducing agent: 2. As a reducing Agent Hydrogen shows greater affinity for oxygen and reduces many metal oxides into free metals. CuO(s) +H2(g) Cu(s) +3H2O(g) WO3(s) + 3H2(g) W(s) + 3H2O(g) CHEMISTRY HYDROGEN AND WATER CHAPTER IX ELEVENTH
  • 59. 59 3. Hydrogenation Reaction: The addition of hydrogen into other molecular compounds called hydrogenation reaction. When molecular compounds and Hydrogen are heated in the presence of Pt or Pd or Ni and other crystals they give addition products. ZnO/Cr2O3 1. CO(g) + 2H2(g) CH3 – OH(l) 400o C /high presusure Methyl Alcohol Ni 2. CH2 = CH2(s) + H2(g) CH3 – CH3 Ethene 400o C Ethane Ni 3. Edible oils(liquid)+ H2(s) Vegetable Ghee (Solid) Un saturated High Temp. Saturated 4. Reaction with metals: Alkali metals like Na, K etc. and alkaline earth metals like Ca, Ba react with hydrogen on heating to form ionic hydrides. 200 o C 1. 2Na(s) + H2(g) 2Na+ H – (s) Sodium Hydride 200 o C 2. Ca(s) + H2(g) Ca+2 H2(s) – 3. Calcium hydride 5. Reaction with non-metals: Hydrogen reacts with many non-metals under different conditions to form addition products. 500 o C/200dtm 1. N2(g) + 3H2(g) 2NH3 (g) Fe2O3/K2O Sunlight 2. H2(g) +Cl2(g) 2HCl(g) 450 o C 3. H2(g) +S(s) H2S(g)
  • 60. 60 Q - 2) What is nascent hydrogen? Describe its reactivity: Nascent Hydrogen: Hydrogen at a time of its birth is chemically more reactive than molecular hydrogen because it is available in atomic form and known as Nascent Hydrogen or Newly born hydrogen. Reactivity: When a piece of Zn metal adds in the acidic Ferric chloride (FeCl3) solution, nascent hydrogen is generated which reduces FeCl3 into ferrous chloride (FeCl2) which is greenish in colour. FeCl3(aq) + H2(g) No reaction Zn/HCl FeCl3(aq) + [H] FeCl2(aq) + HCl(aq) Nascent Hydrogen Ferrous Chloride Acidic KMnO4 (pink) solution can be reduced by nascent hydrogen to colorless solution. Nascent Hydrogen Zn/H2SO4 2KMnO44(aq) + 3H2SO4 + 10[H] K2SO4(aq) + 2MnSO4(aq) +8H2O Q - 3) Define isotope and discus various types of hydrogen isotopes: Isotope: Isotope is the same atom with same number of protons but different number of neutrons in its respective nuclei. Isotopes of Hydrogen: There are three isotopes of hydrogen namely Protium, Deuterium and Tritium. 1. Protium or Ordinary Hydrogen Atom H1 1 : This isotope contains one proton and one electron in the first orbit.  Its atomic number is 1 and mass number is also 1.  About 99.98% of free hydrogen contains Protium.  It is stable isotope of hydrogen. 2. Deuterium (D or H1 2 ): This isotope contains one proton, one neutron and one electron in the first orbit.  Its atomic number is 1 and mass number is 2.  About 0.0156% of free hydrogen contains Deuterium.  It is also stable and heavy isotope of hydrogen i.e. misnomer. 3. Tritium (T or H1 3 ): This isotope contains one proton, two neutrons and one electron in the first orbit.  Its atomic number is 1 and mass number is 3.  About 4 x 10–15 % of free hydrogen contains Tritium.  It is radioactive isotope with half-life of about 12.5 years.  It is used as tracer in the nuclear reactions.
  • 61. 61 Q - 4) How is hydrogen prepared commercially: Industrial Preparation of Hydrogen: Hydrogen is prepared by the following methods. (i) By Passing steam over coke (ii) From Natural Gas (iii) By Thermal decomposition of methane (iv) By the electrolysis of water (i) By passing steam over coke: (Coke-steam process) When steam passes over red-hot coke at about 1000 o C, a mixture of carbon monoxide (CO) and hydrogen (H2) (Water Gas) produced. 1000 o C C(s) + H2O(g) CO(g) + H2(g)⏟ 𝑊𝑎𝑡𝑒𝑟 𝐺𝑎𝑠 Water gas is a very good fuel and used in preparation of methanol (MethylAlcohol).Hydrogen (H2) is separated by two methods from water gas: (i) By liquefaction (ii) By Oxidation a) Separation of H2 by Liquefaction: When water is cooled up to – 200 o C, carbon monoxide liquefies and leaving behind H2 Gas. If the traces of CO gas left in the mixture, than the remaining mixture treats with caustic soda (NaOH) to form sodium formate and leaving behind pure H2 gas. CO(g) + NaOH(aq) HCOONa(aq) Carbon monoxide Caustic Soda Solution 𝑆𝑜𝑑𝑖𝑢𝑚 𝐹𝑜𝑟𝑚𝑎𝑡𝑒 Q. Give Bosch method to separate hydrogen gas from water gas b) Separation of H2 by Oxidation (Bosch Process): This is the most suitable method for the separation of H2 gas from water gas, in this process more steam passes through water gas at 500 o C in the presence of iron oxide (FeO) or chromium oxide (Cr2O3) catalyst. Carbon monoxide gas in water gas oxidizes to CO2 gas, which is soluble in water under pressure, liberating H2 gas. FeO CO(g) + H2(g)⏟ + H2O(g) CO2(g) + 2H2(g) Water gas Water 500 oC Solouble in water
  • 62. 62 (ii) By Natural Gas: (Hydrocarbon-steam process) When steam passes over hydrocarbon such as methane (CH4) which is major constituent of natural gas in the presence of Nickel at about 900 o C, a mixture of carbon monoxide (CO) and hydrogen (H2) (Water Gas) produced. 900 o C CH4(g) + H2O(g) CO(g) + 3H2(g)⏟ Ni 𝑊𝑎𝑡𝑒𝑟 𝐺𝑎𝑠 (iii) By Thermal Decomposition of Methane: Hydrogen Gas (H2) is also commercially prepared by thermal decomposition of methane (CH4), when methane (CH4) is heated in the absence of Air at about 700 o C, methane (CH4) decomposes thermally to produce carbon black C(s) and H2 gas. Above 700 oC CH4(g) C(s) + 2H2(g) Absence of air Carbon Black Uses of Carbon Black C(S):  It is used in rubber industry as filler for manufacturing motor tyers.  It is also used in the preparation of ink, paints, polishers, carbon papers and plastics. (iv) By The Electrolysis of Water: Hydrogen Gas (H2) can also be produced by the electrolysis of water. When electric current is passed through water in the presence of a few drops of acid or base, H2 gas is liberated at cathode and O2 gas a by – product collects at anode. The presence of acid or base helps in the ionization of water. Electricity H2O(l) 2H2(g) + O2(g) Acid or Base Hydrogen gas Q - 5) Give reaction H2 with: (i) Ethene (ii) Ca metal (iii) S(iv) Cl2 (i) Reaction with Ethene CH2 = CH2(s) + H2(g) CH3 – CH3 (ii) Reaction with Ca metal Ca(s) + H2(g) Ca+2 H2(s) – (iii) Reaction with S H2(g) +S(s) H2S(g) (iv) Reaction with Cl2 H2(g) +Cl2(g) 2HCl(g)
  • 63. 63 Q - 6) Define Water and describe some main physical and chemical properties of Water: Water: Water is the most abundant compound on Earth's surface, covering about 70 percent of the planet. In nature, water exists in liquid, solid, and gaseous states. Physical Properties of Water: 1. It is a colorless, odorless and tasteless liquid and excellent solvent. 2. It melts and boils at much higher temperature than other liquids. 3. At 4 o C its density is maximum about 1 g/cm3 . 4. It freezes at 0 °C and boils at 100 °C. 5. It is one of the few substances that expand upon freezing. 6. Water does not obey law of expansion. Chemical Properties of Water: 1. Reaction with metals: (More Electropositive) Alkali metals like Na, K etc. and alkaline earth metals like Ca, Ba react with cold water to form their hydroxides with the liberation of H2 gas. Cold Sodium Hydroxide 1. 2Na(s) + H2O(l) 2NaOH(aq) + H2(g) Cold Calcium Hydroxide 2. Ca(s) + 2H2O(l) Ca(OH)2(aq) + H2(g) (slow) Cold Potassium Hydroxide 3. 2K(s) + H2O(l) 2KOH(aq) + H2(g) Reaction with metals: (Less Electropositive) Less electropositive metals like Mg, Zn or Fe etc. reacts with hot water to liberate H2 gas with formation of their oxides. Iron reacts with excess of steam at red heat. Hot 1. Mg(s) + H2O(l) MgO(s) + H2(g) Hot 2. Zn(s) + H2O(l) ZnO(s) + H2(g) Red Heat 3. 3Fe(s) + 4H2O(g) Fe3O4(s) + 4H2(g) Noble metals like copper, gold, silver and mercury do not react with water in any form.
  • 64. 64 2. Reaction with non-metals: a) With Chlorine Chlorine reacts with water to produce HCl and hypochlorous acid (HClO). Hypochlorous acid is unstable and readily liberates atomic oxygen, which can bleach dyes and kill bacteria by oxidation. Chlorine, therefore in water is both bleaching as well as oxidizing agents. Cl2(s) + H2O(l) ⇌ HCl(aq) + HClO (aq) HClO(aq) 2HCl(aq) + [ O ] O +O O2(g) b) With Carbon When steam is passed over heated coke at 1000 o C, a mixture of hydrogen and carbon monoxide, known as water gas is produced. 1000 o C C(s) + H2O(g) CO(g) + H2(g)⏟ 𝑊𝑎𝑡𝑒𝑟 𝐺𝑎𝑠 c) With Silicon Silicon reacts with steam at very high temperature to form an oxide of silicon i.e. silicon dioxide with liberation of H2 gas. High temp. Si(s) + 2H2O(g) SiO2 (g) + 2H2(g) 3. Reaction with Calcium Oxide (Quick Lime): Calcium oxide partially dissolves in water to form calcium hydroxide i.e. slaked lime. CaO(s) + H2O(l) Ca(OH)2(s) Quick lime Slaked lime 4. Reaction with Nitrous Oxide: Nitrous oxide dissolves in water to form a neutral solution. N2O(g) + H2O(l) Dissolves to form a neutral solution (Nitrous oxide) Q - 7) What do you understand by the anomalous behavior of water? And what is the significant of this unusual behavior of water? Anomalous behavior of Water: water does not obey the law of expansion and contraction between 0 o C to 4 o C and shows anomalous or unusual behavior because water is polar molecule and due to greater polarity all water molecules are associated by means of hydrogen bonding. It expands when cooled form 4 to 0 o C and contracts from 0 to 4 o C. Due to this ice (Solid Water) floats over water and aquatic animals in winter seasons in the region where temperature reaches below 0 o C.
  • 65. 65 Q - 8) What is water of crystallization? What happens when hydrates are heated? Define heat of hydration & write formulas of hydrates. Water of Crystallization: Water of crystallization is the water that contains few molecules of water as a part of the crystal lattice. Crystal Salts containing water of crystallization are called hydrates. When hydrates are heated: When hydrates are heated, the molecules of water of crystallization are easily dissociated form their salts crystals. The residue left behind is then said to be anhydrous (shapeless) or anhydrate. CuSO4.5H2O(s) CuSO4(s) + 5H2O(g) Blue crystals evaporation Anhydrous (white powder) Heat of Hydration: The minimum amount of heat liberated in the formation of hydrate is called heat of hydration. Formulas of Hydrates: 1. Copper Sulphate (CuSO4. 5H2O) 2. Ferrous Sulphate (FeSO4. 7H2O) 3. Sodium Carbonate (Na2CO3. 10H2O) 4. Aluminium chloride (AlCl3. 6H2O) 5. Barium Chloride (BaCl2. 2H2O) 6. Potash Alum [K2SO4. Al2(SO4)3. 24H2O] Q - 9) What is “Potable water”? Write four main characteristics of it. Potable water: The water that is fit for drinking purpose is called potable water. Characteristics of Portable water: 1. It should be free form germs, bacteria and all sorts of pollutants. 2. It should be moderately soft and its hardness should be under 150 ppm (parts per million) 3. It should be free from corrosive substances. 4. It should be colour less odorless and tasteless and have pH rage of 7 – 8.5. Q - 10) What do you mean by soft water, hard water and heavy water? Soft water: The water that containing dissolved impurities but in small quantities and easily produces lather with soap is known as soft water. Hardwater:The water that containing dissolved impurities of hydrogen carbonates, chlorides and Sulphate of calcium and magnesium is called hard water. It does not give lather with soap. Heavywater: The water that containing compound of oxygen with heavy hydrogen i.e. deuterium is known as soft water. Its formula is D2O. Uses of Heavy water: 1. It is used as moderator in nuclear fission reactions to slow down the neutrons. 2. It is used as tracer in biological and chemical researches.
  • 66. 66 Characteristics/Properties of Heavy water: 1. Its density is slightly greater than ordinary water and is 1.104g/cm3 . 2. It has low vapor pressure than ordinary water. 3. Its melting point is 3.81 0 C and boiling point 101.42 0 C. 4. The molecular mass of heavy water is 20 a.m.u. Q - 11) What do you mean by hard water and describes the types of hardness? Also write causes of hardness. Hard water: The water that containing dissolved impurities of hydrogen carbonates, chlorides and Sulphate of calcium and magnesium is called hard water. It does not give lather with soap. Types of Hardness: There are two types of hardness in water. 1. Temporary hardness 2. Permanent hardness Temporary Hardness: Temporary hardness is due to the presence of dissolved hydrogen carbonates of calcium and magnesium. These salts are water soluble and ionize into water as: Ca(HCO3)2(aq) ⇌ Ca2+ (aq) + 2HCO– 3(aq) Mg(HCO3)2(aq) ⇌ Ma2+ (aq) + 2HCO– 3(aq) Permanent Hardness: Permanent hardness is due to dissolved chlorides and Sulphates of calcium and magnesium. These salts are also water soluble and ionize into water as: CaCl2(aq) ⇌ Ca2+ (aq) + 2Cl– (aq) Mg(SO4)(aq) ⇌ Ma2+ (aq) + SO4 2– 3(aq) Causes of Hardness: 1. Presence of bicarbonates of chlorides or Sulphates of calcium or magnesium in water is the main causes of hardness of water. 2. Water hardness is caused by the presence of calcium Sulphate (Ca2+ ) and magnesium hydroxide (Mg2+ ) ions, when they occur in very high concentrations. Q - 12) How is the hardness of water removed? Describe various methods to remove hardness of water? Method to remove hardness of water: The hardness of water can be removed by the following methods. 1. Temporary hardness 2. Permanent hardness a) By heating b) By Clark’s method 4. By Ion – exchange method 5. By using washing Soda 6. By using caustic Soda 7. By using zeolite or permutit
  • 67. 67 Removing Temporary Hardness a) By Heating: Temporary hardness is due to the presence of dissolved hydrogen carbonates of calcium and magnesium which decompose on heating to CaCO3 and MgCO3 which are insoluble in water and easily remove by filtration. Ca(HCO3)2(aq) 𝐵𝑜𝑖𝑙 → CaCO3 ( ↓ 𝑠)(insoluble) + CO2(g) + H2O(l) Mg(HCO3)2(aq) 𝐵𝑜𝑖𝑙 → MgCO3 ( ↓ 𝑠) (insoluble) + CO2(g) + H2O(l) Once Ca2+ ions or Mg2+ ions are out of water, any soap when added to water becomes soft, the water and lather formation occurs. b) By Clark’s method: Temporary hardness can also be by using Clark’s method. In this method temporary hard water containing hydrogen carbonates of Ca and Mg is treated with slaked lime in the tanks to convert them into their insoluble carbonates. These insoluble carbonates settle down at the bottom of tanks while soft water is drained off for the use. Ca(HCO3)2(aq) + Ca(OH)2aq 2CaCO3 (insoluble) + 2H2O(l) Mg(HCO3)2(aq) + Ca(OH)2aq MgCO3 (insoluble) + 2CaCO3 (insoluble) + 2H2O(l) Removing Permanent Hardness a) By Ion – exchange method: In this method in which calcium and magnesium ion from water are removed as insoluble precipitates. The chemicals employed are mostly soluble sodium compounds. These form insoluble precipitates of Ca and Mg ions. b) By Using washing soda: (Na2CO3 10H2O) When washing soda is added to permanent hard water, insoluble CaCO3 and MCO3 are precipitated from the soluble salts of Ca and Mg. CaSO4 (aq) + Na2CO3(aq) CaCO3(s) (insoluble) + Na2SO4(aq) MgCl2 (aq) + Na2CO3(aq) MaCO3(s) (insoluble) + 2NaCl (aq) c) By Using Caustic Soda: (NaOH) When caustic soda is added to permanent hard water, insoluble hydroxide of Mg2+ ion is precipitated from the salts of Ca(OH)2 is partially soluble in water. MgSO4 (aq) + 2NaOH (aq) Ma(OH)2(s) (insoluble) + Na2SO4(aq)
  • 68. 68 By Using Zeolite or Permutit: This is a modern method employed for the softening of hard water. Hydrated sodium Aluminium silicate (Na2Al2Si2O8.xH2O) is called permutit. These complex salts are also known as zeolites. The permutit as loosely packed in a big tank over a layer of coarse sand. Hard water is introduced into the tank from the top. Water reaches the bottom of the tank and then slowly rises through the permutit layer in the tank and hardness is removed. The cat ions present in hard water are exchanged for sodium ions. Therefore this method is also called ion exchange method. Where Z = Al2Si2O8.nH2O CaSO4 (aq) + Na2Z(zeolite) CaZ(s) (insoluble) + Na2SO4(aq) Sodium zeolite can be regenerated by passing a strong NaCl solution through Ca – Zeolite. CaZ (zeolite) + 2NaCl(aq) Na2Z(s) (Sod – zeolite) + CaCl2(aq) Q - 13) Describe disadvantages of hard water and uses of water? Disadvantages of hard water: 1. It consumes more soap in washing process. 2. It may cause stomach disorder or dysentery whenever drinks. 3. If hard water is used in steam engine and turbine then more fuel is used for their heating. 4. If hard water is used in boiler then it reduces the strength of the boiler metal. Uses of water 1. Water is vital for life. Humans, plants and animals cannot survive without water. 2. It is required for irrigating crops, as seeds cannot germinate without water. 3. Water is used in cooking and washing. Running water is used to generate electricity. 4. Water serves as a medium for transportation, as ships and boats move on water. 5. Many industries such as petroleum, fertilizer, dye and drugs industries require large quantities of water for various processes. 6. Fish, other animals and many plants live in water. Fishing and other water- related activities are a source of livelihood for many people.
  • 69. 69 Q - 14) Describe the classification of water Pollutants: water borne diseases that are caused by microorganism presents in water, name various types of water pollutants and their different categories. ClassificationofwaterPollutants:There are various types of water pollutants which can broadly be classified as (a) Oxygen demanding wastes (b) Synthetic organic compounds (c) disease – causing wastes (microorganisms) (d) agricultural water pollutants. a) Oxygen demanding wastes: (OD) These include domestic and animal sewage, bio - degradable organic compounds and industrial wastes from food – processing plants, meat packing plants, slaughter houses, paper and pulp mills, tanneries etc. all these wastes undergo degradation and decomposition due to which there is a rapid depletion of demand oxygen that is harm for aquatic animals. b) Synthetic organic compounds: These are the man – made materials such as synthetic pesticides, synthetic detergents, food additives, pharmaceuticals insecticides, paints, fibers, solvents, plastics etc. these materials are potentially toxic to plants, animals and humans. They cause offensive colors, odors and tastes in water. c) Disease causing wastes: (Microorganisms) These include pathogenic microorganisms which may enter water along with sewage and other wastes and may cause tremendous damage to public health. These microbes comprising mainly of viruses and bacteria can cause dangerous water – borne diseases such as typhoid, cholera, polio, dysentery, and infections hepatitis in humans. d) Agriculture water pollutants: In modern agriculture, pesticides, fertilizers and organic wastes (manure) are essential for producing high yields of crops required for the worlds’ growing population. Some common pesticides used in Pakistan are alder in, DDT, dielderin etc. Q - 15) Describe hygroscopic substances: Hygroscopic Substances: The substances which absorb moisture on exposure to atmosphere are called Hygroscopic Substances. CHARACTERISTICS OF HYGROSCOPIC SUBSTANCES In Solid state they do not form solution but merely become sticky or moist when exposed to atmosphere. Such as sodium nitrate (NaNO3), copper oxide (CuO), Quick lime (CaO). In Liquid state they absorb water from the atmosphere usually diluting itself up to about three times of its original volume. Such as Sulphuric Acid (H2SO4). Uses: They are commonly used in laboratory as drying agents.
  • 70. 70 Q - 16) Name only some common treatment to make municipal water fit for drinking purposes. The names of some common treatment to make municipal water fit for drinking are as follows: (a) Aeration (b) Settling (c) Coagulation (d) Filtration (f) Chlorination Q - 17) What happens when? (i) Mg metal is reacted with hot water. Ans: When Mg metal is reacted with hot water it produces MgO and liberates H2 gas. Hot Mg(s) + H2O(l) MgO(s) + H2(g) (ii) Methane is heated above 700 0 C in the absence of air. Ans: When Methane is heated above 700 0 C in the absence of air it produces Carbon black [C(s)] and liberates H2 gas. Above 700 o C CH4(g) C(s) + 2H2(g) Absence of air Carbon Black (iii) Water gas is heated under pressure in the presence of ZnO – Cr2O3 catalyst. Ans: When water gas is heated under pressure in the presence of ZnO – Cr2O3 catalyst it produces methyl alcohol. ZnO/Cr2O3 CO(g) + 2H2(g) CH3 – OH(l) Water gas 400o C /high presusure Methyl Alcohol (iv) A piece of Zn metal is added to the acidic solution of FeCl3. Ans: When Zn metal is added to the acidic solution of FeCl3 it produces ferrous chloride (FeCl2) and hydrochloric acid (HCl). Zn FeCl3(aq) + [H] FeCl2(aq) + HCl(aq) Nascent Hydrogen Ferrous Chloride
  • 71. 71 Q - 1) Describe Carbon and its uses. Also write its physical and chemical properties: Carbon: Carbon is a pure non – metal and the sixteenth most abundant element in the earth crust. Carbon belongs to IV – A group its atomic number is 6 and atomic weight is 12. Carbon exists in three crystalline i.e. graphite, diamond and Bucky balls. In the Free State carbon occurs in following forms: 1. Crystalline forms: a. Diamond b. Graphite c. Bucky balls 2. Amorphous forms: a. Lamp black b. Wood charcoal c. Coal d. Animal charcoal e. Gas Carbon Physical Properties of Carbon: 1. It is odorless and tasteless solid. 2. All the different carbon allotropes are black or greyish black solids except diamond. 3. It has high melting point above 3000 0 C. 4. It is insoluble in all common solvents like water, alcohols, acids and petrol etc. That’s why Carbon deposits formed during incomplete combustion of fuels (Petrol) inside motor engines have to be removed mechanically; this process is called decarbonizing of motor engines. Uses of Carbon: As Diamond: 1. It is used as gems and precious stones because of its sparkling brilliance. 2. Black diamond is used in drillings, in making of instruments for cutting glasses and metals. 3. Its tiny fragments are used as abrasive for polishing tools. As Graphite: 1. It is used as lubricant to reduce friction in machines, bicycle chains and bearings. 2. Its lined crucibles are used for making high grade steel and other alloys. 3. It is also used in lead pencils and as black pigment in paints. 4. It is also used as neutron moderator in nuclear reactions 5. It is used for making inert electrodes in dry cells and in industrial electrolytic processes. CHEMISTRY CARBON, SILICON AND THEIR COMPOUNDS CHAPTER IX TWELVE
  • 72. 72 As Coal and Coke: 1. Coal and Coke are important fuels and source of energy for homes and industries. 2. It is also used in electric power generator. 3. Coke is used in the extraction of metals from their ores in manufacture of iron and steel. As Charcoal: 1. It is mainly used as a domestic fuel and also as absorbent. 2. Activated charcoal is used in gas masks for absorbing poisonous gases. 3. It is also used for decolorizing of petroleum Gel. As Carbon black: 1. It is used in manufacturing of rubber tyers as filler to increase the strength and hardness of rubber. 2. It is also used in black shoes polishes, printer’s ink, type – writing papers etc. As Carbon fiber: 1. It is used in manufacturing of incorporated plastics to produce a very light but stiff and strong material. Chemical Properties of Carbon: 1. Combustion of Carbon (C) All forms of carbon burn in excess of air (O2) to produce carbon dioxide gas (CO2). Its bond dissociation energy is – 394 K.J/mol. C(s) +O2(g) CO2(g) H = – 394 KJ/mol. In the limited supply of air, incomplete combustion may take place to produce carbon monoxide (CO) instead of CO2 gas. C(s) +O2(g) 2CO(g) 2. Combination reactions: Carbon combines directly with other elements such as hydrogen, Sulphur, calcium, Aluminium at very high temperature to form addition product. C(s) +2H2(g) CH4 (g) (Methane) C(s) +2S2(s) CS2(l) (Carbon disulfide) 2C(s) +Ca(s) CaC2(s) (calcium carbide) 3C(s) +4Al(s) Al4C3(s) (Aluminium carbide) 3. As a reducing Agent Carbon also shows greater affinity for oxygen and reduces many metal oxides into free metals. 1. 2CuO(s) +C(s) ℎ𝑖𝑔ℎ 𝑡𝑒𝑚𝑝 → 2Cu(s) +CO2(g) 3. 2ZnO(s) +C(s) ℎ𝑖𝑔ℎ 𝑡𝑒𝑚𝑝 → 2Zn(s) +CO2(g) 2. Fe2O3(s) + 3C(s) ℎ𝑖𝑔ℎ 𝑡𝑒𝑚𝑝 → 2Fe(s) + 3CO(g) 4. 2PbO(s)+ C(s) ℎ𝑖𝑔ℎ 𝑡𝑒𝑚𝑝 → 2Pb(s) +CO2(g)
  • 73. 73 4. Reaction with strong Oxidizing Agents: Carbon reacts with strong oxidizing agents like hot concentrated nitric acid and cons. Sulphuric acid to liberate CO2 gas. 1. C(s) +4HNO3 (g) (conc) ℎ𝑜𝑡 → CO2(g) + 4NO2(g) + 2H2O(l) 2. C(s) + 2H2SO4 (g) (conc) ℎ𝑜𝑡 → CO2(g) + 2SO2(g) + 2H2O(l) Q - 2) Define allotropy and allotropes. Discuss chief allotropic forms of carbon, also write their uses & structure. Allotropy: The existence of two or more different form of the same element in the same state is called Allotropy. Allotropes: Allotropes are different forms of the same element in the same state and they have same chemical properties but have different physical properties due to different structures or arrangements of the atom. There are three solid allotropic forms of carbon: Allotropic forms of Carbon: 1. Diamond 2. Graphite 3. Bucky balls 1. Diamond: It is one of the crystalline forms of carbon. It is found chiefly in South Africa, Brazil, Australia and India. It may be of blue, green, yellow red or black color. Black color diamonds are called Bort or carbonado. STRUCTURE OF DIAMOND: In diamond each carbon atoms is covalently bonded with four other carbon atoms to give basic tetrahedral unit and forming a giant three dimensional molecule. The C – C bond length is 1.54A0 and bond energy is 347 kj/mol. Physical Properties of Diamond: 1. It is transparent and bright in pure state also known as hardest natural substance. 2. Its refractive index is 2.45µ which is very high. 3. It is bad conductor of electricity. 4. It has very high melting point about 3500 0 C. 5. Its density is about 3.51 g/cm3 . Uses of Diamond: 1. It is used as gems and precious stones because of its sparkling brilliance. 2. Black diamond is used in drillings, in making of instruments for cutting glasses and metals. 3. Its tiny fragments are used as abrasive for polishing tools.
  • 74. 74 2. Graphite: It occurs naturally as Plumbago an opaque black solid. It is found in Siberia, Canada and Sri Lanka. STRUCTURE OF GRAPHITE: In graphite the carbon atoms form flat layers. Each carbon atom in graphite is linked covalently to three other carbon atoms in the same layer to give basic hexagonal ring. Each layer of carbon atoms viewed as two-dimensional sheet polymer or layer lattice at a distance of 1.42 A0 . The distance between the parallel layers is 3.35 0 A. Physical Properties of Graphite: 1. It is dark grey color crystalline solid with dull metallic luster. 2. It is soft and greasy to feel and leaves black mark on paper. 3. It is less dense than diamond. 4. It has density about 2.2 g/cm3. 5. It is good conductor of electricity. 6. It and has high melting point about 3000 0 C. Uses of Graphite: 1. It is used as lubricant to reduce friction in machines, bicycle chains and bearings. 2. Its lined crucibles are used for making high grade steel and other alloys. 3. It is also used in lead pencils and as black pigment in paints. 4. It is also used as neutron moderator in nuclear reactions 5. It is used for making inert electrodes in dry cells and in industrial electrolytic processes. 3. Bucky Balls: is a spherical fullerene molecule with the formula C60. It has a cage- like fused-ring structure (Truncated icosahedron) which resembles a soccer ball, made of twenty hexagons and twelve pentagons, with a carbon atom at each vertex of each polygon and a bond along each polygon edge. Physical Properties of Bucky Balls: 1. They are resilient to impact and deformation. 2. They are not very reactive due to the stability of the graphite-like bonds. 3. They are also fairly insoluble in many solvents and also conduct electricity. Uses of Bucky balls: 1. The antioxidant properties of Bucky balls may be able to fight the deterioration of motor function due to multiple sclerosis. 2. Bucky balls may be used to store hydrogen, possibly as a fuel tank for fuel cell powered cars. 3. Bucky balls may be able to reduce the growth of bacteria in pipes and membranes in water systems.
  • 75. 75 Q - 3) Discus amorphous forms of Carbon. Or Short note Amorphous For Carbon: The amorphous forms of carbon are not considered as allotropes of carbon because x-rays analysis have revealed that they have structures like graphite with the exception of coal which is mined directly from natural deposits. The other amorphous forms can be prepared in various ways. Amorphous Forms of Carbon: (i) Coal (ii) Coke (iii) Charcoal Coal: It is originates from the vegetation of the carboniferous era. It is said that the decomposition of plants and trees occurred gradually under the earth in the absence of air and under pressure CO2 water and methane were liberated leaving behind a material containing high percentage of carbon. During this process under the earth the vegetable material was converted in stages into peat, lignite (Brown Coal), bituminous coal (Soft) and finally to anthracite (Hard Coal). Coal can be mined at various depths from the earth surface and mainly used as fuel. Coke: it is prefaced by heating bituminous coal to very high temperature about 1300 0 C in the absence of air to remove all the volatile constituents present in coal. The process is called destructive distillation of coal. The other non – volatile products obtained by this process would be Coal-Tar and Coke. It burns in air with no smoke and leaves very little residue. Coke is used as fuel and also as reducing agent in the extraction of metals especially iron. Charcoal: It can be produced by heating wood, nut shells, bones, sugar etc. Wood charcoal is most common and prepared by burning wood in the limited supply of air. It may contain impurities such as Sulphur. It is mainly used as domestic fuel. Animal charcoal is produced when animal bones and refuse are heated in the limited supply of air it contains high percentage of calcium phosphate Ca3 (PO4)2 as impurity. It is used sugar industries to remove the brown color from cane sugar and also in decolorizing petroleum Gel. Q - 4) Define catenation: Catenation: It is the ability of the atoms of carbon to bond itself forming long chains and rings and also to from compounds chain and ring together. This is the unique properties of carbon. | | | — c — c — c — | | |
  • 76. 76 Q - 5) What is silicon? How does silicon occur in nature? Silicon: It is a metalloid. It belongs to IV-A group in the periodic table. It is the second most abundant element found in the earth’s crust after oxygen. Its atomic number is 14 and atomic weight is 28. OCCURRENCE OF SILICON: 1. Silicon does not occur in Free states, although silicon is widely distributed in nature. 2. In the combined stated it occurs mainly as silicon IV oxide; SiO2 (Silica) which is present in various forms. Such as sand, quartz, flint, kieselguhr, agate, etc. 3. Silicon occurs as complex silicates with metallic oxides, like Al2O3, CaO, MgO, K2O etc. Q - 6) What are silicates? And Describe some common silicates and their chemical formulas and uses. Silicates: Silicates are those compounds which have a silicon-oxygen anion chemically combined with such metals as Aluminium, calcium, magnesium, iron, potassium, sodium and others to form silicate salts. SOME COMMON SILICATES: Name Chemical Formulas Uses 1. Feldspar K2O.Al2O3.6SiO2 or KAlSi3O8 Ceramics, Glass, Pottery and Abrasive 2. Kaolin (China Clay) Al2O3.SiO2.2H2O Hydrated Crockery 3. Mica K2O3Al2O3.6SiO2.2H2O Or KAl3Si2O10 Hydrated Electrical insulator resistance to high temp. 4. Talc (Soapstone) 3MgO.4SiO2.H2O Hydrated Ceramics 5. Asbestos CaO.3MgO.4SiO2 Or CaMgSi4O12 Heat insulation, Fire – proofing There are more than 1000 silicates present in the earth’s crust. Kaolin and china clay consist of hydrated Aluminium silicates.
  • 77. 77 Q - 7) Give some physical properties of silicon. Physical properties of Silicon: 1. Amorphous silicon is brown colored hygroscopic powder, having specific gravity 2.35. 2. Crystalline silicon is grey in color, opaque lustrous and octahedral crystalline solid with specific gravity 2.49. 3. It is nonvolatile solid with very high melting point about 1410 0C & boiling about 2600 0C. 4. It is hard enough to scratch glass. It is brittle in nature. 5. It is insoluble in most of the common solvents like water but it dissolves in hydrofluoric acid (HF). 6. It is poor conductor of electricity but sometimes used as semi – conductor. Q - 8) Describe chemical preparation of Silicon. 1. By heating a mixture of pure dry sand with Mg. Silicon is also prepared by heating a mixture of pure dry sand (SiO2) & magnesium metal in a fire clay crucible in the absence of air. SiO2(s) +2Mg(s) ℎ𝑒𝑎𝑡 → Si(s) + 2MgO(s) Dilute hydrochloric acid is then added in the reaction mixture to dissolve unreacted Mg metal and MgO formed. The residue left behind contains amorphous silicon. Mg(s) + 2HCl(aq) ℎ𝑒𝑎𝑡 → MgCl2(aq) + H2(g) MgO(s) + 2HCl(aq) ℎ𝑒𝑎𝑡 → MgCl2(aq) + H2O(l) If unreacted SiO2 is left, it can also be removed by dissolving it in hydro-fluoric acid (HF). SiO2(s) + 4HF ℎ𝑒𝑎𝑡 → SiF4(aq) + 2H2O (Silicon tetra fluoride) 2. By passing vapors of SICl4 over heated sodium or potassium metal. WhenthevaporsofSiCl4 arepassedoverheatedsodiumorpotassiummetalinaninertatmosphere, silicon is produced by the reduction process. 1. SiCl4(g) +4Na(s) ℎ𝑒𝑎𝑡 → Si(s) + 2NaCl(s) 2. SiCl4(g) +4K(s) ℎ𝑒𝑎𝑡 → Si(s) + 4KCl(s) 3. By heating SiO2 with coke Silicon is also prepared by heating SiO2 with coke in an electric furnace. This is an industrial method. ℎ𝑒𝑎𝑡 → SiCl4(g) +2C(s) Si(s) + 2CO(g) Eclectic furnace Crystals
  • 78. 78 Q - 9) Give some uses f silicon: Uses of Silicon: 1. Silicon is used in bronze and steel alloys to increase their tensile strength. 2. Very pure silicon is used in making semi – conductors which are of great importance in computers, transistors etc. 3. It is also used for making silicones which are rubber like liquids. 4. It is also used as lubricant, water – repellent, electric insulators. 5. It is also used in paints, varnishes and polishes. 6. Silicone is also used in the preparation of refractory material such as crucible, fire – bricks etc. Q - 10) Describe preparation of Silica and its properties and uses: Silica: Silica occurs naturally in three main crystalline forms namely quartz, tridymite and crysto balite. The commonest of three is quartz. Physical properties of Silica: 1. In pure state Silica exists in a colorless crystalline form. 2. It is macromolecular compound with silicon and oxygen atoms. 3. It is non – volatile and hard. 4. Its melting point is about 15000C. Preparation of Silica: 1. By heating silicon in Air/Oxygen Silica is prepared by heating silicon in air or oxygen. Si (s) +O2(g) ℎ𝑒𝑎𝑡 → SiO2(s) It can also be prepared in hydrated form as a gelatinous precipitate by warming sodium silicate (Na2SiO3) with conc. HCl solution. Na2SiO3(aq) +2HCl (conc) ℎ𝑒𝑎𝑡 → SiO2.H2O + 2NaCl(aq) Hydrated Silica Uses of Silica: 1. It is widely used in making mortar, cement, concrete, glass and refractory silica bricks. 2. Fused silica (quartz glass) is used in making optical lenses and prisms, heat – resisting articles etc. 3. Large quartz crystals are used for lenses of optical instruments. 4. Powered quartz is used in the making of silicon carbide (SiC), Silicon tetra fluoride (SiF4), sodium silicate (Na2SiO3) and silica bricks for lining furnaces. 5. Kieselguhr (SiO2) absorbs liquids readily and is used as absorbent of nitroglycerine (explosive) in the making of dynamite. 6. It is also used in medicines for making dry antiseptic dressings.
  • 79. 79 Q - 11) Describe preparation of Sodium Silicate or water glass and its properties and uses: Sodium Silicate or Water Glass: Sodium silicate is the common name for a compound sodium Meta silicate, Na2SiO3, also known as water glass. Or Sodium silicate dissolves in hot water under pressure to form a viscous liquid which is known as water glass because it looks like ordinary glass. Physical properties of Sodium Silica: 1. It is colorless glass-like solid. 2. Its melting point is about 10900C. 3. Its density is 2.40 g/cm3 . 4. It is very soluble in water. Preparation of Silica: 2. By heating silica and sodium carbonate Sodium silicate is prepared by heating strongly two parts by mass of silica i.e. sand (SiO2) with one part by mass of sodium carbonate (Na2CO3) until the mixture melt. Na2CO3(s) +SiO2(s) 𝑠𝑡𝑟𝑜𝑛𝑔 ℎ𝑒𝑎𝑡𝑖𝑛𝑔 → Na2SiO3(s) + CO2 ↑ (g) It is obtained glass-like solid with melting point 10900C. Uses of water glass: 1. It is used to get silica Gel. 2. It is used as filler in soap industries. 3. It is used for sizing of paper, for fire-proofing of wood and textiles and for making glue. 4. It is also used for making chemical garden. Q - 12) What is Silica Gel? Also write some uses of it. Silica Gel: When an acid is added in a solution of water glass. It turns into a Gel like substance known as Gel. The formula for Gel is SiO2.nH2O. On complete dehydration by heating a hard porous material is obtained known as Silica Gel. Or It is a granular, vitreous, porous form of silicon dioxide (SiO2) made synthetically from sodium silicate (water glass). Silica gel is hard and more solid than common household gels like gelatin or agar. Uses of Silica gel: 1. It is used as good absorbent and in the refining of petroleum. 2. It is used to prevent medicines being spoiled. 3. It is also used to recover valuable vapors form industrial effluents.
  • 80. 80 Q - 1) Describe Nitrogen and its occurrence & uses. Also write its physical and chemical properties: NITROGEN: Carbon is the most common gas preset in the atmosphere and the tenth most abundant element in the earth crust. Nitrogen belongs to V – A group its atomic number is 7 and atomic weight is 14. Nitrogen was discovered by Scottish Botanist, Daniel Rutherford in 1772. It does not sustain combustion. OCCURRENCE OF NITROGEN: 1. Nitrogen occurs in the Free State as N2 gas in air up to 78% by volume and 75% by mass. 2. In combine state nitrogen occurs abundantly in the earth’s crust as nitrates of sodium, calcium and potassium. 3. In combine state nitrogen is found in organic matter such as proteins, urea and vitamin B compounds. In Free State nitrogen in air is important because it dilutes oxygen. Physical Properties of Nitrogen: 1. It is a colorless, odorless and tasteless gas. 2. Pure nitrogen is slightly soluble in water. 3. It is slightly lighter than air. 4. Its boiling point is -196 0 C, while melting points – 210 0 C. Uses of Nitrogen: 1. In the form of nitrous oxide it is used as an anesthetic. 2. Cryopreservation also uses Nitrogen to conserve egg, blood, sperm and other biological specimens. 3. The CPUs in computers use Nitrogen gas to keep them from heating up. 4. It also serves as an oxidation reaction catalyst. 5. Apart from being an oxidizing agent, it can also be used as a flour bleaching agent and rocket fuel. Chemical Properties of Nitrogen: Molecular nitrogen is unreactive because of strong triple bond between two nitrogen atoms(N ≡ N). Its dissociation energy is 941 Kj/mol. 1. Reaction with Hydrogen (H) At very high temperatures & pressures, nitrogen combines directly with H forms ammonia (NH3) N2(s) +3H2(g) 450 0C /200−250 atm → NH3(g) 2. Reaction with Oxygen (O) At 2000 0 C temperatures, nitrogen combines directly with oxygen and forms nitric oxide NO(g). N2(s) +O2(g) 2NO(g) 3. As a reducing Agent At very high temperatures, nitrogen combines directly with magnesium and forms magnesium nitride Mg3N2 (g). N2(s) +3Mg2(s) Mg3N2(g) CHEMISTRY NITROGEN AND OXYGEN CHAPTER IX THIRTEEN
  • 81. 81 Q - 2) Describe Oxygen and its occurrence & uses. Also write its physical and chemical properties: OXYGEN: Oxygen is the most abundant element in the earth. Oxygen belongs to V–A group its atomic number is 6 and atomic weight is 12. Oxygen was discovered by Scheele in 1772 and Priestley in 1774. However the major properties of oxygen is given by Lavoisier. It is most essential substance for all living things. OCCURRENCE OF OXYGEN: 1. Oxygen occurs in the Free State as well as combined state. 2. In Free State oxygen is present as diatomic gas (O2) in the earth’s atmospheric air up to 21% by volume and about 33% by mass. 3. In combine state oxygen accounts for nearly 50% by mass of the earth’s crust. Physical Properties of Oxygen: 1. It is a colorless, odorless and tasteless gas. 2. It is neutral to moist litmus paper. 3. It is slightly soluble in water only about 2% by volume at room temperature. 4. Gaseous oxygen is about 1.1 times denser than air. 5. It liquefies at -183 0 C and solidifies at – 210 0 C. Uses of Oxygen: 1. This gas is used in various industrial chemical applications. 2. It is used to make acids, sulfuric acid, nitric acid and other compounds. 3. Hot oxygen air is required to make steel and iron in blast furnaces. 4. Some mining companies use it to destroy rocks. 5. Oxygen gas is used to destroy bacteria. Chemical Properties of Oxygen: Oxygen reacts with metal, nonmetals and other compounds directly. 1. Reaction with Calcium (Ca) Oxygen combines directly with Calcium forms Calcium oxides (CaO) 2Ca2(s) +O2(g) heat → 2CaO(s) 2. Reaction with Lithium (Li) Oxygen combines directly with Lithium forms Lithium oxides (Li2O) 4Li(s) +O2(g) heat → 2Li2O(s) 3. Reaction with Ferric sulphide (FeS) Oxygen combines directly with Ferric sulphide forms Ferric oxides (Fe2O3) & Sulphur dioxide. 4FeS2(s) +7O2(g) heat → 2Fe2O3(s) + 4SO2(g)
  • 82. 82 4. Reaction with Sulphur (S) Oxygen combines directly with Sulphur forms Sulphur dioxides (SO2) S(s) +O2(g) iginitio 𝑛 → SO2(g) 5. Reaction with Carbon (C) Oxygen combines directly with Carbon forms Carbon dioxides (CO2) C2(s) +O2(g) iginitio 𝑛 → CO2(g) 6. Reaction with methane (CH4) Oxygen combines directly with Methane forms Carbon dioxides and water. CH4(g) +2O2(g) combussion → CO2(g) + 2H2O(g) 7. Reaction with Hydrogen sulphide (H2S) Oxygen combines directly with Hydrogen sulphide forms Sulphur dioxides and water. 2H2S(g) +3O2(g) combussion → 2SO2(g) + 2H2O(g) Q - 13) Describe chemical preparation of Nitrogen. How can you get nitrogen from the atmospheric air? Give tow uses of nitrogen. 1. From Air The only important method of producing nitrogen gas is the fractional distillation of liquid air. In this process air is firs liquefied to form liquid air which is ten fractionally distilled. Air liquefied by successive compression and expansion. Fractional distillation of liquid air 1. Clean air is compressed and then cooled by refrigeration, upon expanding the air further cools and liquefies. 2. The liquid air is filtered to remove carbon dioxide solid and then distilled. 3. Nitrogen is the most volatile component, with boiling point -196 o C, distill over. 2. In laboratory Pure nitrogen in the laboratory is prepared by heating ammonium nitrate which thermally decompose to give nitrogen gas. Ammonium nitrite is first obtained by reacting ammonium chloride with sodium nitrite. a. Formation of Ammonium Nitrite NH4Cl(s) +NaNO2(s) heat → NH4NO2(s) + NaCl(s) b. Preparation of N2 NH4NO2(s) heat → N2(g) + 2H2O(l) Uses of Nitrogen: 1. In the form of nitrous oxide it is used as an anesthetic. 2. Cryopreservation also uses Nitrogen to conserve egg, blood, sperm and other biological specimens.
  • 83. 83 Q - 14) Describe laboratory preparation of Oxygen. How oxygen is industrially produced from liquid air? 1. From Air The isolation of oxygen from air involves two steps. (a) Liquefaction of air (b) Fractional distillation of liquid air. a. Liquefaction of air 1. Air in the gaseous form is first passed through caustic soda to remove CO2 present in air. 2. It is then compressed under very high pressure about 200 atm in the compressor. 3. It is then cooled and allowed to expand rapidly through a nozzle. 4. The process of compression and expansion are repeated over and over again due to which temperature falls up to – 200 0 C at which air liquefies. b. Fractional distillation of liquid air 1. The liquid air is then led to a fractionating column through a filter in order to remove the traces of CO2 solid if left behind. 2. On distillation nitrogen with lower boiling point of -190 0C, evolves first leaving behind a liquid very rich in oxygen. 3. On heating turns liquid argon into gas which boils out at – 185.7 0 C and passes off from the middle of the column and liquid oxygen. 4. The least volatile component in the air turns into oxygen gas at -183 0 C. 2. In laboratory Oxygen in the laboratory is prepared by heating potassium chlorate mixed with little manganese dioxide, which acts as a catalyst, the decomposition reaction takes place at lower temperature and at much faster rate. 2KClO3(s) MnO2|heat → 2KCl(s) + 3O2(g) Q - 15) What are oxides how are they classified describe normal oxides in detail? OXIDES: The binary compounds of oxygen with metals and nonmetals are called as oxides. Such as CaO, Fe2O3, CO2, H2O etc. Classification of oxides On the basis of valence number or oxidation of oxygen, oxides are classified into several groups. (a) Normal Oxides (b) Peroxides, (c) Super oxides (d) Sub oxides a. Normal oxides Normal oxides are those oxides in which oxygen shows normal oxidation state or valence number -2. It is further divided into four types. (i) Basic oxides (ii) Acidic oxides (iii) Amphoteric oxides (iv) Neutral oxides
  • 84. 84 (i) Basic oxides The normal oxides of metals are the examples of basic oxides. 1. 2Ca2(s) + O2(g) → 2CaO(s) 2. 2Pb(s) + O2(g) → 2PbO(s) 3. 4Na(s) + O2(g) → 2Na2O(s) Most of these are soluble in water and produce their hydroxides. They also turns red litmus to blue. 1. CaO(s) + H2O(l) → Ca(OH)2(aq) 2. Na2O(s) + H2O(l) → 2NaOH(aq) They also react with acids to form salts and water. 3. MgO(s) + 2HCl(aq) → MgCl2(aq) + H2O(l) 4. CaO(s) + 2HNO3(aq) → Ca(NO3)2(aq) + H2O(l) (ii) Acidic oxides The normal oxides of non-metals are generally acidic. 1. S(s) + O2(g) → SO2(g) 2. C(s) + O2(g) → CO2(g) 3. N2(s) + 2O2(g) → 2NO2(g) These oxides react with water to form acids which turns blue litmus to red. 5. SO2(g) + H2O(l) → H2SO3(aq) (Sulphurous acid) 6. Na2O5(s) +H2O(l) → 2HNO3(aq) (Nitric acid) 7. CO2(s) +H2O(l) → H2CO3(aq) (Carbonic acid) They react with alkalis to form salts and water. 8. CO2(g) + 2NaOH(aq) → Na2CO3(aq) + H2O(l) 9. SO3(g) + 2KOH(aq) → K2SO4(aq) + H2O(l) (iii) Amphoteric Oxides The oxides that possess dual characteristics i.e acidic as well as basics are known as amphoteric oxides. 1. 4Al(s) + 3O2(g) → 2Al2SO3(s) 2. 2Zn(s) + O2(g) → 2ZnO (s) They react with alkalis (acids/bases) to form salts and water. With acids Al2O3(s) +6HCl(aq) → 2Al2Cl3(aq) + 3H2O(l) ZnO(s) +H2SO4(aq) → ZnSO4(s) + H2O(l) With bases Al2O3(s) +2NaOH(aq) → 2NaAlO2(aq) (sod aluminate) + H2O(l) ZnO(s) +2NaOH(aq) → NaZnO2(s) (sod zincate)+ H2O(l) (iv) Neutral Oxides The oxides that are neither acidic nor basic are known as neutral oxides. They are neutral to litmus in aqueous solutions. Example: 1. Nitric oxide (NO) 2. Carbon monoxide (CO) 3. Water (H2O) 4. Nitrous Oxide (N2O)
  • 85. 85 b. Peroxides Peroxides are oxides containing higher proportion of oxygen as compared to normal oxides. In these oxides  Oxygen has oxidation state or valence number 1.  They contain peroxide ion (O – O)  They produce hydrogen peroxide with acids. Example: 1. Sodium peroxide (Na2O2) 2. Barium Peroxide (BaO2) Na2O2(s) +2HCl(aq) NaCl(aq) + H2O2(aq) (Hydrogen peroxide) c. Super oxides Super oxides are oxides containing higher proportion of oxygen as compared to peroxides. In these oxides  Oxygen has oxidation state or valence number −0.5 or − 1 2 .  They show tendency to release oxygen (O2) on heating and powerful oxidizing agent.  They do not produce hydrogen peroxide with acids. Example: 1. Potassium superoxide (KO2) 2. Cesium superoxide (CsO2) 3. Rubidium superoxide (RbO2) d. Sub oxides Sub oxides are oxides containing less quantity of oxygen than normal oxides. In these oxides  They are unstable.  Very few sub oxides are known. Example: Carbon sub oxide (C3O2) Q - 16) Give the preparation and properties of hydrogen peroxide and its uses: Hydrogenperoxide(H2O2): The Nard was the first to prepare hydrogen peroxide by the action of dilute Sulphuric acid on barium peroxide (BaO2). He discovered that hydrogen peroxide contain one more oxygen atom in its molecule than water thus called it as oxygenated water. H2SO4(aq) + BaO2(s) Heat → BaSO4(s) + H2O2(aq) Physical Properties of Hydrogen Peroxide: 1. Pure Hydrogen peroxide is a pale blue syrupy liquid. 2. It mixes with water to give solution which is slightly acidic. 3. Its boiling point is 150 o C but it boils with decomposition. 4. It freezes at about – 0.9 o C.
  • 86. 86 Preparation of Hydrogen Peroxide: Hydrogen peroxide is prepared by two Methods. (i) Laboratory Method (ii) Industrial Method (i) Laboratory Method: The hydrogen peroxide is usually prepared in laboratory by the action of dilute Sulphuric acid on peroxides of certain metals, especially barium peroxide (BaO2). Barium peroxide is insoluble and can be easily removed by filtration and pure H2O2 is obtain. H2SO4(aq) + BaO2(s) Heat → BaSO4(s) + H2O2(aq) ← ( 𝐻𝑦𝑑𝑟𝑜𝑔𝑒𝑛 𝑃𝑒𝑟𝑜𝑥𝑖𝑑𝑒) (ii) In Industrial Method/Preparation: The hydrogen peroxide is usually prepared in laboratory by the action of dilute Sulphuric acid on peroxides of certain metals, especially barium peroxide (BaO2). Barium peroxide is insoluble and can be easily removed by filtration and pure H2O2 is obtain. CH3 − CH | OH − CH3(l) + O2(g) → H2O2(l) + CH3 − O || C − CH3(l) Chemical Properties of H2O2: 1. As showing exothermic reaction When hydrogen peroxide is exposed to air, it decomposes to form water and oxygen. The decomposition is exothermic. 2H2O2(l) exposure → 2H2O(l) + O2(g) + ∆H 2. As Oxidizing agent Hydrogen peroxide is common oxidizing agent,usually in the form of aqueous solution with 3% H2O2. It is strong oxidizing agent because it can readily donate oxygen or accept electrons. H2O2(l) → H2O(l) + O(g) (donation of Oxygen) H2O2(l) + 2H+ + 2e– → H2O(l) (acceptor of electron) 3. As reducing agent Hydrogen peroxide in the form of aqueous solution with 3% H2O2 can also behave as a reducing agent. It reduces acidic potassium permanganate solution recoloring of KMnO4. 2KMnO4(aq) 3H2SO4(aq)+ 5H2O2(l) → K2SO4(aq)+ 2MnSO4(aq) +H2O(l) + 5O2(g) It reduces chlorine to hydrochloric acid and Oxygen gas is given off. H2O2(aq) + Cl2 → 2HCl+ O2(l) Isopropyl alcohol Hydrogen peroxide Acetone
  • 87. 87 Uses of Hydrogen peroxide (H2O2): 1. It is used as a mild antiseptic in mouth wash as well as for cleaning wounds. 2. It is used as bleaching agent in bleaching delicate materials like silk, wool, feathers and human hairs. 3. It removes unwanted color from fabrics, hair or other materials. 4. Its liquid is used for restoring painting & providing oxygen for burning fuel in space rocket. 5. It is used in the preparation of compounds like sodium chlorate (NaClO2). Q - 17) Define oxidation and reduction and write chemical equation in support of each. Oxidation: The process or reaction in which oxygen combines with other elements or compounds to produce oxides. It is also known as addition of oxygen. Reduction: The process or reaction in which removal of oxygen occurs from a substance. It is also known as removal of oxygen. Oxidation involves (a) Addition of Oxygen (b) Removal of Hydrogen (c) Loss of electrons a. As addition of oxygen When oxygen reacts with iron, magnesium or carbon produces their oxides. 4Fe(s) + 3O2(g) → 2Fe2O3(s) (Ferric oxide) Mg(s) + O2(g) → 2MgO (s) (Magnesium oxide) C(s) + O2(g) → CO2(g) (carbon dioxide) b. As removal of hydrogen By this process removal of hydrogen occurs form a compound. H2S(g) + Cl2(g) → S(s) + 2HCl c. As loss of electron By this process loss of electron occurs from a substance. Sn(s) → Sn2+ + 2e – Reduction involves (a) Addition of Hydrogen (b) Removal of Oxygen (c) Gain of electrons a. As addition of Hydrogen By this process Hydrogen is added to a substance. H2S(g) + Cl2(g) → S(s) + 2HCl b. As removal of Oxygen By this process removal of oxygen occurs form a compound. CuO(s) + H2(g) → Cu(s) + H2O(l) c. As gain of electron By this process gain of electron occurs from a substance. Al3+ + 3e– → Al(s)
  • 88. 88 Q - 18) Define oxidizing agent, reducing agent and redox reactions. Oxidizing Agent: A substance that accepts or gains electrons is known as oxidizing agent. And a substance itself reduced. Reducing Agent: A substance that donates or loss electrons is known as reducing agent. And a substance itself oxidized. Redox Reaction: The reaction in which oxidation & reduction occur simultaneously is called oxidation-reduction reaction or redox reactions. Q - 19) What is ozone? How ozone is produced in the atmosphere? Or how oxygen is converted into ozone? Its physical/chemical properties and what is important of ozone and write uses of ozone. OZONE: Ozone is pale-blue poisonous gas with a sharp, irritating odor. It is an allotropic form of oxygen with molecular formula O3. It was first discovered by Schonbein in 1839. OCCURRENCE OF OZONE: 1. It exists in a layer at a height of about 20 kilometers above the earth. 2. Very small amount of ozone is produced around electrical machineries when they are in operations. Physical Properties of Ozone: 1. It is a bluish, colored gas that has a boiling point of -112 oC. 2. It is very poisonous gas at concentration 100 parts per million (ppm). 3. It is only slightly soluble in water but dissolves in turpentine oil readily. 4. It has characteristic smell which is sharp irritating like Cl2 gas. Uses of Ozone: 1. It is used in treatment of domestic water in place of chlorine. 2. It is used as bleaching because all oxidizing agents are also good bleaching agent. 3. It is largely used in the preparation of pharmaceuticals, synthetic lubricants and other commercially useful organic compounds. Important of Ozone: It protects the earth from the harmful effects of high energy rays. But In lower atmosphere ozone is measured as air pollutant it damages living system. Chemical Properties of Ozone: Ozone is chemically more reactive than ordinary diatomic oxygen. It acts as powerful oxidizing agent because Ozone dissociates readily forming reactive oxygen atoms. O3 → O2 + O(g) ∆H = −107 KJ/Mol.
  • 89. 89 1. Reaction with Lead sulphide (PbS) Ozone oxidizes lead sulphide (PbS) in acidic medium liberating oxygen (O2) gas. PbS2(s) +4O3(g) → PbSO4(s) + 4O2(g) 2. Reaction with Sulphuric Acid (H2SO4) Ozone oxidizes Hydrogen sulphide (H2SO4) in acidic medium liberating oxygen (O2) gas. H2SO4(s) +2O3(g) → H2SO4(aq) + O2(g) 3. Reaction with Sulphur dioxide (SO2) Ozone oxidizes Sulphur dioxide (SO2) in acidic medium liberating oxygen (O2) gas. SO2(s) +O3(g) → SO3(g) + O2(g) 4. Reaction with Potassium Iodide (KI) Ozone oxidizes Potassium Iodide (KI) in acidic medium liberating oxygen (O2) gas. KI(aq) +O3(g) + H2SO4(s) → K2SO4(aq) + I2(aq) + H2O(l)+ O2(g) Preparation of Ozone (O3): Ozone is prepared by two Methods. (i) Atmospheric Method (ii) By electric discharge Method (i) Atmospheric Method: In nature ozone is formed form atmospheric oxygen by lightning flashes however ozone is very unstable dissociates readily forming reactive. O3 → O2 + O(g) ∆H = −107 KJ/Mol. (ii) By Electric discharge method: Ozone can be prepared from oxygen by passing electric discharge through oxygen gas. It is necessary to use silent discharge because sparking would generate heat energy which decomposes ozone proceed. The apparatus used for converting oxygen into ozone is known as Ozonizer. 3O2(g) Electric discharge → 2O3(g) Q - 20) What is Aqua Regia? How does it dissolve gold? Aqua Regia: The mixture of concentrated nitric acid (HNO3) and hydrochloric acid HCl, optimally in a volume ratio of 1:3 is called Aqua Regia which is also known as "Royal Water". Aqua Regia dissolves gold due to liberation of nascent chlorine which forms gold chloride with it, which is soluble. HNO3(conc) + 3HCl(conc) → NOCl + 2H2O(l) + 2Cl NOCl → NO + Cl Au(s) + 3Cl → AuCl3 Oxygen Ozone Nitrosyl chloride Nascent chlorine Nitrosyl chloride
  • 90. 90 Q - 21) Give the preparation and properties of Ammonia (NH3) and its uses: Ammonia (NH3): Ammonia is a very important chemical in industry, in nature ammonia is produced during the decay of nitrogenous matter in the absence of air. Physical Properties of Ammonia: 1. It is colorless gas with a characteristic pungent smell. 2. It is highly soluble in water about 1300ml dissolves in 1ml of water. 3. It is easily liquefied into a colorless liquid at ordinary temperature by compression. 4. In large quantity it is poisonous because of its effects on respiratory system. Preparation of Ammonia: Ammonia is prepared by two Methods. (i) Laboratory Method (ii) Industrial Method (i) Laboratory Method: In the laboratory ammonia is prepared by heating ammonium salts usually ammonium chloride (NH4Cl) with slaked lime i.e. calcium hydroxide. 2NH4Cl(s) + Ca(OH)2(s) Heat → CaCl2(g) + 2H2O(l) + 2NH3(g) ← (Ammonia) (ii) In Industrial Method/Preparation: (Haber – Bosch process) On large scale ammonia is manufactured by the direct combination of Nitrogen and Hydrogen from Haber – Bosch process. In this process a mixture of pure Nitrogen and Hydrogen in the ratio of 1:3 by volumes is allowed to react. The basic problem in ammonia synthesis is that it is a reversible reaction and can be described as To get maximum yield of ammonia  The Optimum temperature should be 400 – 450 oC.  The pressure should be 200 – 250 atm.  The suitable catalyst Fe2O3 (Ferric oxide) with small amount of Al2O3, CaO K2O.