Chemistry finisher

730 views
649 views

Published on

Published in: Technology
0 Comments
0 Likes
Statistics
Notes
  • Be the first to comment

  • Be the first to like this

No Downloads
Views
Total views
730
On SlideShare
0
From Embeds
0
Number of Embeds
0
Actions
Shares
0
Downloads
19
Comments
0
Likes
0
Embeds 0
No embeds

No notes for slide

Chemistry finisher

  1. 1. By: Johanna Louise B. Neri
  2. 2. Chapter 1. The Nature of ChemistryKey Concepts Chemistry - is the study of matter and the changes it undergoes. It is sometimes called the central science because it overlaps with many other sciences. Technology - is the application of science. It has improved the quality of human life. SI units - are used to express physical quantities in all sciences. Metric prefixes are used to make units smaller or larger.
  3. 3.  Precision - is how close several measurements are to the same value. Accuracy - tells how close a measurement is to the true or accepted digit. Significant - figures include both the certain digits and the estimated digit. Scientific notation - is used to write very small or very large numbers.
  4. 4.  Dimensional analysis - is the technique that uses conversion factors. The guide to ensure that conversion factor are properly formulated is the cancellation of units.
  5. 5. Chapter 2. Matter: Its Compositionand Organization. Key Concepts Matter - is anything that has mass and volume. Properties of matter differ for solids, liquids, and gases. A pure substance is either an element or a compound. An element is a substance that cannot be broken down to simpler substances. A compound is formed when two or more elements combine in a chemical change.
  6. 6.  A change in the properties of substance without a change in composition is a physical change. If there is a change in the composition of a substance, a chemical change has occurred. Chemical changes produce matter with new properties. The physical combination of two or more substances is a mixture. A mixture has a variable composition. It may be heterogeneous or homogeneous. Heterogeneous mixtures (coarse mixtures, suspensions, and colloids) do not have uniform properties throughout, while homogeneous mixtures (solutions) have uniform properties.
  7. 7.  Solutions may be gases, liquids, or solids. The components of a mixture can be separate by physical methods. Colloids - are mixtures of two or more solids, liquids, or gases whose particles are bigger than the particles of a solution but smaller than those of a suspension. Tyndall effect, Brownian movement, adsorption, and electrical charge are the properties of colloids.
  8. 8.  Colloids are prepared and purified by condensation and dispersion methods. Condensation - is the process of combining molecules to form colloidal particles. Dispersion - is the process of breaking down large particles to colloidal size. Energy - is the capacity to do work or to transfer heat. It is involved whenever matter undergoes a change.
  9. 9. Chapter 3. Atomic Theory.Key Concepts Over 2400 years ago, the concept of the atom was proposed by Greek philosophers. In the early 19th century, Daltons proposed the atomic theory. This theory is related to the three fundamental laws of matter. (1) The total mass of the reactants and products are constant during a chemical reaction (law of conservation of mass).
  10. 10.  (2) Any sample of compound, has elements in the same proportion (law of definite composition). (3) In different compounds of the same elements, the mass of one element that combines with a fixed mass of the other can be expressed as a ratio of small whole numbers (law of multiple proportions).
  11. 11.  Thomson’s experiment on the behavior of cathode rays in magnetic and electric field led to the discovery of the electron and the measurement of its charge to mass ratio. Millikan’s oil drop experiment measured the charge of the electron. Becquerel and the Curies discovered radioactivity. Rutherford’s studies on alpha rays led to the discovery of nucleus.
  12. 12.  Atoms have a nucleus that contains protons and neutrons. Electrons move in the space around the nucleus. Elements can be classified by atomic number or the number of protons in the nucleus of an atom. All atoms of a given element have the same atomic number. The mass number of an atom is the number of protons and neutrons. All atoms of the same element that differ in mass number are known as isotopes.
  13. 13. Chapter 4. Electronic Configuration.Key Concepts The properties of visible light and other forms of electromagnetic radiation led to the electronic structure of atoms. Max Planck proposed that energy is absorbed and emitted in discrete amounts or individual packets called quanta (plural for quantum). Albert Einstein used Planck’s theory to explain the photoelectric effect. He proposed that light consists of quanta of energy which behave like tiny particles of light. He called these energy quanta photons.
  14. 14.  The concept of quantized electrons grew from the study of line spectra of atoms. A line spectrum consists of quanta of energy which can be used like fingerprints to identify the element. Niels Bohr used the line spectra to explain specific energy levels within the atom. He proposed the planetary model of the atom. Louis de Broglie discovered the wave nature of matter which initiated the development of a new mathematical description of electron configuration.
  15. 15.  Heisenberg’s uncertainty principle explained the impossibility of simultaneously measuring the momentum and location of an electron. Erwin Schrodinger devised the quantum mechanical model of the atom which described electrons as waves that exist in quantized energy levels. The regions in space around the nucleus where electrons are most likely to be found are called orbitals. These orbitals have various shapes and are labeled s, p, d, and f. Each principal energy level or shell consists of these orbitals.
  16. 16.  The manner in which electrons are arranged around the nucleus of an atom is called electron configuration. The Aufbau principle, the Puali exclusion principle, and the Hund’s rule are applied in writing electron configurations. The Aufbau principle tells the sequence in which orbitals are filled. The Pauli exclusion principle states that a maximum if only two electrons can occupy an orbital. Hund’s rule explains that electrons pair up only after each orbital on a sublevel is occupied by a single electron.
  17. 17. Chapter 5. The Periodic Table.Key Concepts Different periodic table were developed by Dobereiner, Newlands, Mendeleev, and Meyer. The periodic table was based on similarities in properties and reactivities of elements in the increasing order of their atomic mass. Discrepancies in these periodic tables were resolved when Moseley established that each element has a unique atomic number and showed that elements should be arranged according to their increasing atomic number.
  18. 18.  The periodic table is organized into 18 groups or families and 7 periods or rows. The groups are organized further into s, p, d, and f blocks based on how valence electrons fill each sublevel. Elements in a group have similar properties because they have the same valence electrons. Atomic radius decreases from left to right across a period because the positive charge of the atoms increases, which attracts electrons more strongly.
  19. 19.  Atomic radius increases down a group because the electrons of the atoms fill more energy levels. Ionization energy - is the energy absorbed to remove an electron to form a positive ion. Electron affinity - is the energy when an atom gains an electron forming a negative ion. Electronegativity is the attraction of an atom for electrons in a chemical bond.
  20. 20.  The trends for ionization energy, electron affinity, and electronegativity ate the same. They increase from left to right of the periodic table and decrease down a period. Metals are found on the left side of the periodic table. Nonmetals are found on the upper right side of the periodic table. Metalloids have some properties of metals and nonmetals.
  21. 21. Chapter 6. Chemical Bonds.Key Concepts Chemical bonds are classified into three groups: ions of opposite charges; covalent bonds, which result from the sharing of electrons by two atoms; and metallic bond, which are the attractions among positively charged ions for delocalized electrons. These bonds involve the valence electrons with the tendency of atoms follow the octet rule. This can be represented by electron – dot symbols or Lewis symbols. Resonance structures are used when a simple Lewis structure is not adequate to represent a particular molecule or ion (specie). Some covalent molecules formed from atoms of the representation groups 1, 2, and 3 lack octet configurations while atoms from 5, 6, and 7 form expanded octet configurations.
  22. 22.  A polar covalent bond is formed when electrons are not shared equally between two atoms. Electronegativity difference of bonded atoms determines the kind of bond formed between the atoms. The sharing of one pair of electrons produces a single bond, the sharing of two pairs, a double bond, and three pairs, a triple bond. Double and triple bonds are also called multiple bonds.
  23. 23. Chapter 7. Molecular Geometry.Key Concepts The shapes of small molecules can de explained in terms of the VSEPR model which states that electron pairs arrange themselves as far apart as possible to minimize electrostatic repulsion. The geometry of molecules is determined by the arrangement of bonding pairs and lone pairs. The five common shapes of small molecules are linear, trigonal planar, tetrahedral, trigonal bipyramid, and ictahedral.
  24. 24.  The electron pair cloud repulsion model suggests that the denser the electron clouds, the greater the repulsive force. The order from greatest to least repulsive force is that triple bond > double bond > lone pair > single bond (≡>═>1.p.>─). Molecules that contain polar bonds (bond dipoles) may be polar or nonpolar molecules, depending on the shape of the molecules. The properties of polar molecules (dipole) are different from those of nonpolar molecules. Valence bond theory - is an extension of the Lewis covalent bond. In this theory, bonds are formed when neighboring atoms overlap and the potential energy of the system decreases. The greater the overlap, the stronger the bond formed.
  25. 25.  Shapes of molecules are also described in terms of hybrid orbitals. The process of hybridization involves the promotion of electron to empty orbital(s) and mixing of the orbitals to form equivalent numbers of hybrid orbitals. Hybrid orbitals can overlap with orbitals of other atoms to make bonds. Or they can accommodate lone pairs. Covalent bonds that overlap end to end along the line connecting the atoms are called sigma (σ) bonds. When p orbitals overlap on a side to side orientation perpendicular to the line connecting the atoms, these are called pi (π) bonds.
  26. 26. Chapter 8. Chemical Names and Formulas.Key Concepts The charges or oxidation numbers of the ions of representative elements are determined by their position in the periodic table. Most transition metals have more that one common ionic or oxidation numbers. A polyatomic ion is a group of atoms that behaves as an ion – ide. If cations have more than one ionic charge, a Roman numeral is used in the name. Ternary ionic compounds contain at least one polyatomic ion. The names of these compounds end in – ite or – ate. Binary molecular compounds are composed of two nonmetallic elements. Prefixes are used to indicate the number of atoms each element that are present in a molecule of the compound.
  27. 27.  Binary acid are compounds that contain hydrogen and nonmetal ions. They are named by using the prefix hydro followed by the name of the anion ending in – ic acid. Ternary acid contain hydrogen and polyatomic ions. They are named by using the name of the polyatomic ion ending in – ic or - ous acid Based are compounds containing a metal ion and hydroxide ion(OH‾). Bases are named by writing the name of the cation followed by hydroxide. Salts are named by using the name of the cation followed by the name of the anion.
  28. 28. Chapter 9. Chemical Reactions.Key Concept Chemical reactions are represented by chemical equations. The substances that undergo chemical changes are the reactants and the substances formed are the products. Chemical equations must be balanced to be consistent with the law of conservation of mass. In balancing an equation, appropriate coefficients are placed before the formulas of the reactants and products so that the same number of atoms of each element appears on each side of the equation.
  29. 29.  The state of a substance in an equation is detonated by (s), (1), and (g) for solid, liquid, and gas, respectively. A substance dissolved in water is denoted by (aq) for aqueous. If heat, light, or electricity is used to initiate the reaction, its process or symbol is written above the arrow. If a catalyst is used to increase the speed of reaction, its formula or symbol is also written above the arrow. In a combination reaction, two or more elements or compounds combine to produce a single product.
  30. 30.  In a decomposition reaction, a single compound is broken into two or more simpler substances. In a single replacement reaction, a more chemically active element displaces a substance below it in the activity series. A double replacement reaction involves the exchange of cations and anions between two compounds. Replacement reactions can be written as net ionic equations. In a combustion reaction, oxygen is always one of the reactants.
  31. 31. Chapter 10. Stoichiometry.Key Concept A mole is the amount of substance that contains 6.02 ×1023 particles or species. The representative particles of elements are the atoms. Molecules are representative particles of molecular compounds and diatomic elements. The representative particles for ionic compounds are formula units. The mass of a mole of atoms, molecules, or ions is its formula weight expressed in grams called molar mass. A mole is defined in terms of the number of particles in a substance or the mass in grams of the substance. The mole can be used in converting among different units.
  32. 32.  Percent composition of a compound is the percent by mass of each element in a compound. Empirical formula is the simplest whole-number ratio of atoms of elements in a compound. This can be calculated from the percent composition of a compound. Molecular formula shows the actual number of atoms of each element in a compound. It may be the same as or a multiple of an empirical formula. Stoichiometry is the study of the quantitative relationship of individual compounds in chemical reactions.
  33. 33.  The coefficients in a balanced equation represent the relative number of moles of each substance. Coefficients are used in establishing conversion factors as mole ratios in solving stoichiometric problems. The conversion factor relates the mole of a given substance to the moles of the required substance. Units such as grams and particles are converted to moles when solving stoichiometric problems. When reactants supplied are not in the exact amounts required by the balanced equation, that which is used up is the limiting reagent and that which remains after the reaction is completed is the excess reagent.
  34. 34.  The theoretical yield is the amount of product obtained when all of the limiting reagent is used up. The actual yield is the product formed when the actual reaction is carried out. The percent yield is the ratio of the actual yield to the theoretical yield expressed in percent.
  35. 35. Chapter 11. GasesKey Concept The physical properties of gases are given by four quantities:  Pressure P  Volume V  Temperature T  Amount of Gases n The behavior of gases can be explained by the kinetic molecular theory. The standard temperature and pressure (STP) is 0°C and 1 atm. Atmospheric pressure is the pressure exerted by the gases (air) around us which is 1 atm or 760 mm HG.
  36. 36.  Boyle’s law states that the pressure and volume of a gas are inversely proportional to its absolute temperature (constant n and T). Charles law states that the volume of a gas is directly proportional to its absolute temperature (constant n and P). Avogadros law states that equal volumes of gases contain the same number of particles (constant T and P). Ideal gas equation PV=nRT is a combination of the gas laws. Daltons law states that the pressure of a mixture of gases is the sum of the partial pressure of the component gases. Real gases behave like ideal gases in ordinary conditions except at high pressure and low temperature. Lighter gases diffuse and effuse faster than heavier gases do.
  37. 37. Chapter 12. Liquids and SolidsKey Concept At room temperature, substances with weak intermolecular forces of attraction are gases; those with moderate intermolecular forces are liquids; and those with strong intermolecular forces are solids. Intermolecular forces include ion-dipole forces, dipole-dipole forces, London dispersion forces and hydrogen bonds. Physicals properties of liquids and solids are explained by the kinetic molecular theory.
  38. 38.  Liquids possess properties such as viscosity, surface tension, capillarity evaporation, boiling point, and critical temperature and pressure. Heating curve is a plot of temperature versus heat for phase changes. The properties of solids are explained based on their nature and strength if intermolecular forces of attraction. A phase diagram indicates the states or phases of a substance under specific temperatures and pressures.
  39. 39. Chapter 13. SolutionsKey Concept Solutions are homogeneous mixtures of two or more substances in a single phase. A solutions is made of solute, the substance that dissolves, and solvent, the substance in which the solute is dissolved. A substance that dissolves in another substance is soluble (miscible) and if it does not, it is insoluble (immiscible). Solutions are either gaseous, liquids, or solid solutions. In preparing dilute solutions form concentrated solutions, the number of moles before dilutions is equal to the number of moles after dilutions.
  40. 40.  Saturated solutions contains the maximum amount of solute it can dissolve at a given temperature. Unsaturated a solutions that contains less than the maximum. Supersaturated a solution with more than the maximum. Solubility is the extent to which a solute dissolves in a given solvent.
  41. 41. Chapter 14. Chemical KineticsKey Concept Chemical kinetics is the study of rate and sequence of steps by which chemical reactions occur. The rate of a reaction is the measure of how reactants turn into products. Collisions theory assumes that particles collide at the proper orientation and with sufficient energy in order to react. Activation energy is the minimum energy required for a chemical reaction to occur and make the reactant form an activated complex or transition state.
  42. 42.  The factors that affect the rate at which a chemical reaction proceed are nature of the reactants, concentration of the reactants, temperature at which reaction occurs. A rate law for a reaction describes the relationship between the concentration of reactants and the reaction rate. Most chemical reactions proceed through a series of elementary steps. The series of steps called the reaction mechanism. The slow reaction in a reaction mechanism called the rate-determining step.
  43. 43. Chapter 15. Thermo chemistryKey Concept Thermodynamics is the study of processes which involve heat transfer and the performance of work. Thermochemistry is the study of this heat exchange and work on chemical reactions. Energy + Energy = constant: law of conversation of energy. 3 types of system :  Open  Closed  Isolated An open system allows the transfer of both energy and matter into and out the system through a boundary or wall.
  44. 44.  A closed system is only capable of transferring energy through boundary. An isolated system is not capable of transferring both energy and matter into and out of the system through a boundary or wall. Heat is a transfer of energy between system and surrounding due to temperature difference.
  45. 45. Chapter 16. Chemical EquilibriumKey Concept Equilibrium is a state at which there is “balance of forces”. 3 types of equilibrium:  Mechanical  Thermal  Chemical Chemical equilibrium is achieved when the rate of the forward reaction is equal to the rate of the reverse reaction and the amount of components remains unchanged.
  46. 46.  Reversible reactions is an incomplete reactions. The reaction is represented by using a double headed arrow (═). Law of mass reaction states that the compositions of a reaction mixture can vary according to the quantities of components that are present.
  47. 47. Chapter 17. Acids and BasesKey Concept The operational definitions of acids and bases are based on experimental results from the laboratory which includes color change using dyes. Arrhenius acids is a neutral substance that ionizes when it dissolves in water to give the H+ or hydrogen. Arrhenius base is a neutral substance that gives the OH-, or hydroxide ion when dissolves in water.
  48. 48.  Lewis defines an acids as species that can accept a pair of electrons while a base is a species that can donate a pair of electrons. The degree of ionization, not the concentration, classifies an acid or a base as weak or strong. Compounds with more than one proton to give are called polyprotic acids.
  49. 49. Chapter 18. ElectrochemistryKey Concept Electrochemistry is the branch of chemistry that deals with electricity and its relation to chemical reactions. A chemical reactions were loss of electron(s)is involved id called oxidation while reaction where electron(s) is gained is called reduction. Redox reaction can be balanced by using the oxidation number method or the ion electron method.
  50. 50.  Electrochemical cell, voltaic cell, or galvanic cell converts chemical energy from spontaneous reaction to produce electricity. Electrochemical cell is composed of the electrodes and charge carriers. Anode is the electrode where oxidation occurs. Cathode is where reduction occurs or where electrons are accepted. There 3 types of electrodes:  Inert  Metallic  Membrane
  51. 51. Chapter 19. Nuclear ChemistryKey Concept Many elements have at least one radioactivity isotope or radioisotope. Elements with atomic numbers 83 or greater are all radioactivity. Radioactivity decay of naturally occurring radioisotope produces alpha particles, beta particles, and gamma radiations. The half-life of a radioisotope is the time it takes for one-half of a sample of the isotope decay. In artificial radioactivity or artificial transmutation, the nucleus of an atom is bombarded with a particle or radiation and changed into different nuclei.
  52. 52.  In balancing nuclear equation, the sum of the mass numbers and atomic numbers of reactants must be equal to the sum of the mass numbers and atomic numbers of the product. The mass defect in a nucleus is due to the strong forces of attraction that bind nucleons together.
  53. 53. Chapter 20. Organic ChemistryKey Concept Organic compounds are basically made up of carbon atoms bonded mostly to hydrogen, oxygen, nitrogen, and sulfur. Organic chemistry the study of the carbon-based compounds. Hydrocarbons are made up of carbons and hydrogens. Alkanes also called saturated hydrocarbons, have an sp3 hybridization, four sigma bonds with no pi bonds that can be bound to H or C atoms.
  54. 54.  Alkenes are hydrocarbon containing a carbon-carbon double bond. Alkynes are hydrocarbons containing a carbon-carbon triple bond. Cycloalkanes are aliphatic cyclic (alicyclic) compounds which have general ring structure containing –CH-. A molecule can only be aromatic if it has the following properties:  (1) the molecule is planar and  (2) has a monocyclic system of conjugation with a total of (4n + 2) p electrons where n is an integer.

×