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C20 Review Unit 01 Matter Energy And The Periodic Table

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  • 1. Chemistry 20 – Review Unit 1 Matter Energy and The Periodic Table What is Chemistry? Chemistry deals with: a) what things are made of ( matter ). b) properties of matter ( characteristics ). c) how matter interacts ( chemical reactions ). d) energy changes ( energy released or absorbed ). Important Terms in Chemistry: a) Matter : anything that has mass and takes up space (has  volume ). b) Composition : what kinds of particles are in the matter and how many of each. c) Structure : how the particles are bonded together. d) Properties : characteristics of matter that are used to identify it. e) Transformations : changes in matter *
  • 2. Three types of changes: i) Phase Change : state of matter ( solid, liquid or gas ) Example: H 2 O (s)  H 2 O (l) - Solid – - Liquid – - Gas – - Plasma – definite shape and definite volume varied shape and definite volume varied shape and varied volume gas charged with high energy electricity ( Northern Lights – Fluorescent lights ) ii) Chemical Change : atoms within a substance are rearranged to form new substances Example: 2 H 2 O (l)  2 H 2(g) + O 2(g) iii) Nuclear Change : elements are changed into different elements * Example :
  • 3. Energy : something that gives you the ability to do work Scientific Methods Scientists are always making observations . Observations can be of two types: Qualitative : Quantitative : descriptive statements , using the 5 senses , about what has been observed For example: colour , texture , descriptive size etc measured observations, using instruments For example: 5.0 g (mass), 2.75 cm (length), 2.50 L (volume) Properties of a substance are used to identify a substance. There are 2 types of properties: *
  • 4. Physical Properties :   Chemical Properties : the characteristics of a substance that do not involve the change in internal composition of a substance For example: colour , shape , density , melting or boiling point characteristics of a substance that cause it to change (react) into something totally new For example: a substance’s ability to react with an acid *
  • 5. Classification of Matter Matter Pure Substances Mixtures Elements Compounds Metals Nonmetals Metalloids Noble Gases Ionic Molecular Homogeneous Heterogeneous Solution Acids Bases Salt Solutions Alloys Suspension Colloid Physical Changes Chemical Changes
  • 6. Elements : matter that is made up of one kind of atom Examples : - metals such as copper ( Cu (s) ) or mercury  ( Hg (l) ) ( monatomic ) - nonmetals such as helium ( He (g) ), oxygen ( O 2(g) )( diatomic ) , ozone ( O 3(g) ) ( triatomic ) , iodine  ( I 2(s) ) , sulfur ( S 8(s) ) , phosphorous  ( P 4(s) ) ( polyatomic ) Compounds : substances that are made up of groups of elements chemically bonded and arranged in a specific manner Pure Substances
  • 7. Special Compounds : Minerals : Oxides: an element or compound that occurs naturally in the earth Example : salt ( halite ) a compound formed by combining at least one element with oxygen Example : Fe 3 O 4(s) – rust ( metal + oxygen ) CO 2(g) – carbon dioxide ( nonmetal + oxygen )
  • 8. There are two types of compounds: Ionic Compounds : Molecular Compounds : formed by combining metallic elements with nonmetallic elements Examples : - sodium chloride – NaCl (s) - magnesium hydroxide – Mg(OH) 2(s) - aluminum oxide – Al 2 O 3(s) formed by combining nonmetallic elements into a unit called a molecule Examples: - carbon monoxide – CO (g) - water – H 2 O (l) - ammonia – NH 3(g) - wax – C 25 H 52(s)
  • 9. Mixtures ( Impure Substances ) - contain 2 or more pure substances that are not chemically joined together and mixed in any proportion - can be physically separated by filtering , boiling or centrifuging - Example: a container of rocks , sand and water There are two types of mixtures: Homogeneous : - looks completely uniform in composition - has only one phase - usually are solutions Examples : - sugar dissolved in water (solid in liquid) - air (gas in gas) - carbonated water (gas in liquid)
  • 10. Heterogeneous : - does not look uniform in composition - distinctly visible phases Examples : - oil and vinegar salad dressing - chunky peanut butter Special Mixtures : Colloid : falls halfway between a homogeneous and heterogeneous mixture as it looks homogeneous but distinct phases exist upon closer, microscopic observations Examples : - milk (liquid in liquid) - gelatin (liquid in solid) Ore : a rock that has an element that can be obtained for profit Examples : - iron ore – Fe 2 O 3(s) - bauxite – Al 2 O 3(s)
  • 11. Alloy : a mixture of at least two elements , at least one of which is a metal , melted together uniformly Examples : - steel – iron and carbon - stainless steel – iron, carbon, chromium and manganese Plated Metals : - a metal object that has been coated with another metal Examples : - Nails may be coated with zinc and steel cans are coated with tin to prevent corrosion . - Some dinner utensils are coated with silver to look more attractive .
  • 12. Changes in Matter Physical Changes: - There are three major types of energy changes: - phase, chemical and nuclear . Phase Change: - There is no change in chemical composition , only the change in phase is noted. As a substance changes from one phase to another, it either absorbs energy ( endothermic change ) or releases energy ( exothermic change ). - Example: H 2 O (s)  H 2 O (l)
  • 13. Solid Liquid Gas Plasma Melting Fusion Vapourization Ionization Sublimation De-ionization Condensation Freezing Solidification Deposition
  • 14. Chemical Change: - The reactant(s) have different chemical composition than the product(s). - Evidence of a chemical change occurring: - colour change - formation of a solid ( precipitate ) in a solution - production of a gas ( bubbling ) - energy change ( heat being lost or absorbed ) Example: 2 H 2 O (l)  O 2(g) + 2 H 2(g)
  • 15. Nuclear Change or Transmutation: - One kind of element is changed to form another kind of element . Fission : splitting of atoms Examples: Alpha Decay Beta Decay Fusion : atoms coming together Example: atomic mass protons alpha emission - α: alpha particles or helium nuclei ( ) electron emission
  • 16. Energy and Matter Kinetic Energy : - energy of motion - the mass of an object and how fast it is moving determines the total energy it has Potential Energy : - stored energy - may be due to: - an object’s vertical position relative to another object - energy used to push atoms apart - used to hold subatomic particles together Example : A car travelling at has less energy than when it is moving at .
  • 17. Examples : - a ball held 1 m above the ground has more potential energy than being held 0.5 m above the ground - gasoline has potential energy stored in chemical bonds - splitting atoms releases stored nuclear energy that holds the subatomic particles together
  • 18. Law of Conservation of Energy Energy cannot be created nor destroyed but can only be transformed from one form to another . Energy Kinetic Potential Heat Mechanical Sound Electrical Light Chemical Nuclear Gravitational
  • 19. The Periodic Table - Mendeleev created the periodic table to organize elements together based upon similarities of chemical reactions . - It is made up of 18 vertical columns called groups or families and horizontal rows or series called periods . Period Alkali Metals Alkaline Earths Nonmetals Halogens Noble Gases Lanthanum Series Actinium Series 1 2 18 17 16 15 14 13 3 4 5 6 7 8 9 10 11 12 1A 2A 3A 4A 5A 6A 7A 8A 1 2 3 4 5 6 7 Transition Metals
  • 20. Characteristics of Metals - Shiny or lustrous - Good conductors of heat and electricity - Flexible: can be bent or shaped into wire ( ductile ) or rolled into sheets ( malleable ) - All are solids at room temperature (exception: mercury , which is a liquid ) - All have a silver colour , except for gold and copper Examples: copper ( Cu (s) ), magnesium ( Mg (s) )
  • 21. Characteristics of Nonmetals - Solids may be dull in appearance or may have a glassy luster - Poor conductors of heat and electricity - Varied examples of solids, liquids and gases - Colours vary Examples: oxygen ( O 2(g) ), sulfur ( S 8(s) ), bromine   ( Br 2(l) ) Characteristics of Metalloids - Have properties of both metals and nonmetals - Semiconductors : conduct in some conditions but not others - All are solids Examples: carbon ( C (s) ), boron ( B (s) ), silicon ( Si (s) )
  • 22. Characteristics of Noble Gases - All are gases - Least reactive of all elements - Do not conduct electricity Examples : helium ( He (g) ), neon ( Ne (g) ) - Each element has an abbreviation. Some elements use only one character such as carbon (C) or potassium (P). - Each element has an abbreviation. Some elements use only one character such as carbon (C) or potassium (K). Others use two characters , the first being capitalized and the second , being lower case , such as calcium (Ca). - Some symbols are obvious but some have historic origins.
  • 23. Examples : Common Name Symbol Latin Name antimony Sb stibnum copper Cu cuprum gold Au aurum iron Fe ferrum lead Pb plumbum mercury Hg hydragyrum potassium K kalium silver Ag argentum sodium Na natrium tin Sn stannum
  • 24. The Structure of the Atom Models of the Atom Dalton Model – Billiard Ball Model - Atoms are single , indivisible spheres . - Molecules are units or groups of atoms arranged in a specific ratio . Problem : It was found that a) the atom contained positive and negative parts. b) the positive part was much heavier than the negative part. c) the whole atom was electrically neutral d) opposite charges attract and similar charges repel .
  • 25. Thomson Model – Raisin Bun Model - The atom is a large mass of positive charge and has small negative parts embedded in this sphere. Rutherford Model – Nuclear Model or Empty Space Model - The atom has a central core that is positively charged . - Electrons exist in empty space , surrounding this central positive mass, travelling anywhere they want. 3+
  • 26. Bohr Atom – Orbital Model - The core of the atom was the nucleus , containing neutral particles called neutrons and positive particles called protons . - Electrons had different defined amounts of energy , therefore had to exist at specified distances from the nucleus , in orbits . - This meant that the electron was quantized or had only a certain quantity of energy when found at different energy levels . 3+
  • 27. Quantum Mechanical Model – Cloud Model - The electrons move so quickly in their orbital that they create a “ cloud-like ” behavior. - Each “cloud” has its own characteristic shape , depending how far away from the nucleus the energy level is found. Subatomic Particles and The Structure of the Atom Atom: - the basic building block of all substances - contains a central nucleus which houses relatively similarly - sized neutrons ( no charge ) and protons ( positive charge ) - electrons exist in orbitals around the nucleus Isotopes: - Atoms that have the same number of protons and electrons (therefore are the same element) but have different numbers of neutrons - Some elements may have more than two isotopes
  • 28. Isotopes of Hydrogen Protium (ordinary hydrogen) Deuterium (heavy hydrogen) Tritium (radioactive hydrogen) 1 p + 0 n o 1 e – 99.985 % abundance 1 p + 1 n o 1 e – 0.015 % abundance 1 p + 2 n o 1 e – negligible abundance
  • 29. Particle Symbol Charge Actual Mass (g) Mass Relative to a Proton Mass Relative to an Electron proton p + +1 1.672 x 10 –24 1 1836 neutron n o 0 1.675 x 10 –24 1 (1.002) 1839 electron e – – 1 9.11 x 10 –28 0 (0.0005) 1
  • 30. Atomic Mass - used to be called atomic weight - the atomic mass described on a periodic table is the average mass of all known isotopes of that element - carbon is known to have 6 different isotopes, mass numbers ranging from 10 to 16 - about 98.89 % of all the isotopes have a mass of 12.00000 amu (atomic mass units) and is called carbon – 12 - about 1.11 % of all the isotopes have a mass of 13.00335 amu and is called carbon – 13 - the other four isotopes are very rare and are not found naturally - if the periodic table gives an atomic mass of carbon as 12.01, that means that the average mass of all carbons is 12.01 - all atomic masses on the periodic table have a unit called grams per mole   (we will discuss the concept “mole” later)
  • 31. Electron Energy Level Representations for Atoms - using the concepts from the Bohr Model of an atom we can give a simplistic sketch of how atoms are put together Sodium Atom - since the element sodium has an atomic number 11 on the periodic table, this means that every sodium atom must have 11 protons ( 11 positive charges ) - since atoms are neutral , there must also be 11 electrons - the atomic mass is given as 22.99 ( round off to 23 ) - so, the total mass of a sodium atom is 23 , made up of subatomic particles inside the nucleus ( 11 protons and 12 neutrons – electron mass is too small to count )
  • 32. The Sodium Atom ( Na  ) 1 st energy level = 1 st period 2 nd energy level = 2 nd period 3 rd energy level = 3 rd period 2 e – 8 e – 1 e – 11 p + 12 n o The outside energy level contains 1 electron, therefore, tells you that sodium is found in Group IA. The Fluorine Atom ( F  ) 2 e – 7 e – 9 p + 10 n o The outside energy level contains 7 electrons, therefore, tells you that fluorine is found in Group VIIA. 1 st energy level = 1 st period 2 nd energy level = 2 nd period
  • 33. Ions - positive ions are called cations and negative ions are called anions - ions are atoms that have lost or gained one or more electrons - the electrons are lost or gained from the last , outside energy level only - the electrons on the outside energy level are called valence electrons - if an electron is lost , a negative charge is lost - losing a negative makes the atom more positive - if an electron is gained , a negative charge is gained - gaining a negative makes the atom more negative
  • 34. The Sodium Ion ( Na + ) - since the sodium ion is shown as a +1 charge , that means that it has lost an electron from the outside energy level 11 p + 12 n o 2 e – 8 e – 1 e – 1 st energy level = 1 st period 3 rd energy level = 3 rd period 2 nd energy level = 2 nd period
  • 35. - therefore, the sodium ion now looks like this: - the total charge is +1 11 p + 12 n o 2 e – 8 e – 1 st energy level = 1 st period 2 nd energy level = 2 nd period 11 p + 12 n o 2 e – 8 e – 1 e – 1 st energy level = 1 st period 3 rd energy level = 3 rd period 2 nd energy level = 2 nd period
  • 36. The Fluoride Ion ( F – ) - note that if you look up the names of the ions formed from nonmetals , their names change to an “ ide ” ending - to get an F – ion, you must add one electron to the atom to get a total charge of –1 9 p + 10 n o 2 e – 8 e – 1 st energy level = 1 st period 2 nd energy level = 2 nd period
  • 37. Elements, Compounds and Nomenclature - the term “ nomenclature ” refers to “ naming ” - chemical nomenclature is the organized system chemists use to name substances and write their chemical formulas - when writing any formula for any substance we first assume that the substance exists on its own , at room temperature - if special conditions exist, states of matter are adjusted
  • 38. Naming Pure Elements Metals - when naming metals , we simply state the elemental name given on the periodic table - when writing the formula for any pure metal we simply write the elemental symbol ( no subscripts ) - inclusion of states of matter is very important, so we must always indicate a solid state, the exception being mercury , Hg (l)
  • 39. Nonmetals - when naming nonmetals , we simply state the elemental name given on the periodic table - when writing the formula for any pure nonmetal we write the elemental symbol with subscripts for some of the nonmetals, because we must also be aware of how it exists at room temperature - for most nonmetals, you’ve just got to memorize their formulas and their natural states of matter - all nobles gases are monatomic : He (g) , Ne (g) , Ar (g) , Kr (g) , Xe (g) , Rn (g) - diatomic elements include: - nitrogen – N 2(g) - oxygen – O 2(g) - all halogens : fluorine – F 2(g) chlorine – Cl 2(g) bromine – Br 2(l) iodine – I 2(s) astatine – At 2(s) - polyatomic elements include: phosphorous – P 4(s) sulfur – S 8(s) - all other nonmetals are monatomic and solid
  • 40. Naming Compounds - a compound is a pure substance made by combining at least two different elements in a specific ratio Ionic Compounds - an ionic compound is formed by taking one metallic ion and combining it with one nonmetallic ion or a complex ion Binary Ionic Compounds - only one metallic ion and one nonmetallic ion are combined Process for Creating the Formula: - find the elements - list their ions - place positive ions ( cations ) first , followed by negative ions ( anions ) - criss-cross their charges to create the correct ratio of each element and simplify charges (reduce) - the state of matter for all ionic compounds is solid
  • 41. Naming Ionic Compounds - the first element gets its normal elemental name as found on the periodic table - the second element’s name has an “ ide ” ending (regardless of the ratios found in the formula) Examples: - sodium and chlorine Na + and Cl – join together Na + Cl – criss-cross charges Na –1 Cl +1 simplify subscripts NaCl (one’s are ignored) identify state of matter NaCl (s) name the compound sodium chloride
  • 42. - magnesium and iodine Mg 2+ and I – Mg 2+ I – Mg –1 I 2+ MgI 2 MgI 2(s) magnesium iodide - oxygen and aluminum Al 3+ and O 2– Al 3+ O 2– Al –2 O +3 Al 2 O 3 Al 2 O 3(s) aluminum oxide - oxygen and calcium Ca 2+ and O 2– Ca 2+ O 2– Ca –2 O +2 Ca 2 O 2 *note: subscripts are simplified for ionic compounds CaO CaO (s) calcium oxide
  • 43. Creating the Formula and Naming Ionic Compounds Using Complex Ions - when a simple positive ion is combined with a complex ion the process is as follows: - list both positive and negative ions - combine both - use brackets for the complex ion - criss-cross their charges - simplify subscripts - when naming the compound, use the normal first name for the metallic ion and copy the given name of the complex ion , which is found in the complex ion table Example: - calcium and hydroxide Ca 2+ and OH – Ca 2+ and (OH) – Ca –1 (OH) 2+ Ca 1 (OH) 2 - since 1’s are not required Ca(OH) 2(s) - note that if a 1 should appear after the brackets, the brackets are not required
  • 44. Naming Ionic Compounds Using Multiple Ion Charges - some metallic ions have more than one charge - for example, iron has Fe 2+ and Fe 3+ - in these cases it is necessary to actually state which ion is being used - Roman numerals are used after each multiple-charged ion Example iron and oxygen Fe 2+ and O 2­– FeO (s) named iron (II) oxide or Fe 3+ and O 2­– Fe 2 O 3(s) named iron (III) oxide - some tables may use old, “ classical ” names that end with “ ic ” or “ ous ” - “ic” ending is for the ion that has the greater charge Example Fe 3+ is also called ferric Fe 2+ is also called ferrous FeO (s) is also called ferrous oxide Fe 2 O 3(s) is also called ferric oxide
  • 45. Hydrated Compounds - some compounds have a strange-looking formula which has water added at the end , such as CuSO 4   5 H 2 O (s) - these are still ionic compounds and exist as solids at room temperature - the only thing we have to do is to state how many waters are involved - so, for CuSO 4   5 H 2 O (s) , the first part is named as copper (II) sulfate and then we add that there are 5 waters - the Latin prefix for 5 is penta and water is called hydrate - the name becomes copper (II) sulfate pentahydrate
  • 46. mono di tri tetra penta hexa hepta octa nona deca # Latin Prefix 1 2 3 4 5 6 7 8 9 10
  • 47. Naming Molecular Compounds - molecular compounds are formed by combining nonmetallic elements - when these elements are combined they may form more than one compound - carbon and oxygen may combine to form CO (g) or CO 2(g) - nitrogen and oxygen may combine to form NO (g) or NO 2(g) or N 2 O 4(g) - note that subscripts are not simplified and another naming system is used
  • 48. Example - CO 2(g) is made from carbon and oxygen - the first element gets the normal name - the second element gets and “ ide ” ending - now we must state how many of each element is being used by inserting prefixes before each element name - there is one carbon ( mono carbon ) and two oxygens ( di oxide ) - if the first element is a “mono” the prefix is ignored but is used for the second element - NO (g)­ is nitrogen monoxide - NO 2(g) is nitrogen dioxide - N 2 O 4(g) is dinitrogen tetroxide
  • 49. - there is no general rule for determining states of matter - each molecular compound has its own characteristic state of matter at room temperature and these are learned as you go along - some molecular compounds have classical names that have no logic involved in their naming system - these we just memorize Examples H 2 O (l) is NH 3(g) is O 3(g) is CH 4(g) is C 6 H 12 O 6(s) is C 12 H 22 O 11(s) is CH 3 OH (g) is C 2 H 5 OH (l) is H 2 O 2(l) is H 2 S (g) is water ammonia ozone methane glucose ethanol sucrose methanol hydrogen peroxide hydrogen sulfide
  • 50. Hydrogen Compounds - hydrogen compounds are those which contain a hydrogen at the beginning of the chemical formula and are dissolved in water ( aqueous ) - Examples: HCl (aq) , HNO 3(aq) - hydrogen bonds covalently ( shares electrons ) to nonmetals to form a molecular compound which may be any state of matter , depending on the species being formed - most hydrogen compounds are named as acids - the only exceptions to this rule are the following pure substances - HCl (g) – - HCl (g) – hydrogen chloride - H 2 S (g) – - H 2 S (g) – hydrogen sulfide - HCN (g) – - HCN (g) – hydrogen cyanide - when hydrogen compounds dissolve in water they form acidic solutions - HCl (g) is bubbled into water to form a solution called hydrochloric acid
  • 51. Properties of Acids - turns litmus indicator red - tastes sour - neutralizes bases - conducts an electrical current - pH is lower than 7 Naming Acids Naming acids is easy if we follow this table: hydrogen ___ide becomes hydro___ic acid hydrogen ___ate becomes ________ic acid hydrogen ___ite becomes ________ous acid
  • 52. Examples: Name the following acids: - HF (aq) – the normal name given to this chemical is hydrogen fluor ide , therefore, hydrogen fluoride becomes hydro fluor ic acid - HNO 3(aq) – the normal name given to this chemical is hydrogen nitr ate , therefore, hydrogen nitrate becomes nitr ic acid - HNO 2(aq) – the normal name given to this chemical is hydrogen nitr ite , therefore, hydrogen nitrite becomes nitr ous acid We can read the table backwards to write out the chemical formula of a given acid name. Example : Give the chemical formula for hydrosilicic acid. - the acid name came from hydrogen silic ate , therefore the formula must be H 2 SiO 3(s) - now change states to give the acid formula, H 2 SiO 3(aq)
  • 53. Classification of Acids Binary Acids : Binary Acids : contain a hydrogen and one other kind of atom Example: HCl (aq) Oxo Acids : Oxo Acids : contain a hydrogen , an oxygen and one other kind of atom Example: HNO 3(aq)