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1 Chapter 2 Atoms, Molecules, and Ions 2 The Atomic Theory Proof atoms exist Laws of Chemical Combination Atomic Theory of Matter: Dalton (1808) Hypotheses were derived from 3 Laws of Chemical Combination 1. All matter is composed of tiny particles called atoms 2. All atoms of a given element are identical, but atoms of different elements differ in size, mass and chemical properties F H Fluorine 18 amu 3. Hydogen atom 1 amu Compounds are formed when atoms of different elements unite in fixed proportions. 74amu = 72amu F F 4. F F + 2amu + H H 38amu + 36amu = 74amu F H F F + H F A chemical reaction involves atomic rearrangement. No atoms are created or destroyed 3 4 Law of Definite Proportions A compound will always have same chemical composition Each product formed from definite proportions of reactants Same mass proportions & atomic ratios of elements present 8 Water molecules 16H + 8O 8H2O 5 Law of Multiple Proportions When 2 different compounds made of the same 2 elements are compared, the masses of one element combines with a fixed mass of the second element. The combination is in a ratio of small whole numbers. Open face sandwich 1 bread + 1 filling 1:1 ratio Regular sandwich 2 bread + 1 filling: 2/1 2:1 ratio Compounds are formed when atoms of different elements unite in fixed proportions 6 Law of Conservation of Mass Total mass is constant during a chemical reaction Mass of reactants must exactly match mass of products 2lbs potatoes + 3 ounces milk+ 1 ounce butter = 2.25 lbs mashed potatoes Also called a mass balance 11.1g H2 (g) + 88.9g O2 (g) = 100.0g H2O (l) No matter can be created or destroyed 7 Structure of the Atom What’s on the inside The Electron and Cathode Rays (1890’s) Electricity causes ray to be deflected in a vacuum Few molecules: minimal molecular interference Applied magnetic field or electric field deflects ray Ray is attracted to positively charged anode, so it must be made of negatively charged particles 8 9 Millikan Oil Drop Experiment Finalized related charge and mass of electron JJ Thompson Ratio of cathode ray particle’s charge to mass =–1.76 X 108 C/g Negatively charged particle is called an electron Robert Millikan Charge on an electron e = -1.602 X 10-19 C Mass of an electron me= 9.10 x 10-28g 10 The Atom J. J. Thomson’s Raisin Pudding Model Positively charged sphere with electrons imbedded inside Rutherford’s Experiments Shot particles through gold foil Results: Hit detector: No interference Deflected from straight line: Charge interference Reflected back toward emitter: Hit + small center The Atom and Sub-Atomic Particles Protons large, positively charged particles in small central nucleus Electrons tiny, negatively charged particles in cloud around the nucleus Neutrons large, neutral particle in nucleus electron neutron nucleus proton Elements are not electrically charged Must have equal numbers of protons and electrons 11 12 Neutrons Atomic particle with same mass as proton Discovered by James Chadwick No charge- accounts for extra mass of isotopes Change in mass, no change in chemical properties 13 Atomic Number, Mass Number and Isotopes Differences between atoms 14 Atomic Symbols Atomic Number (Z) # protons in a nucleus Determines element identity Located lower left on symbol Mass Number (A) # protons + # neutrons Determines isotope identity Located upper left on symbol A E Z 35 Cl 17 15 Isotopes Elements with the same number of protons and electrons, but differing number of neutrons Used in chemistry for structure identification or to follow a particular molecule through a reaction Example: hydrogen and deuterium Water, H2O 1 Heavy water, D2O 2 H has 1 proton and 0 neutrons Most abundant D (deuterium) has 1 proton & 1 neutron Occurs 1/6700 molecules H 1 1 H 16 The Periodic Table Elemental Organization 17 Mendeleev’s Periodic Table 1869 Arranged the known elements in order of increasing atomic mass from left to right and from top to bottom in groups. Elements with similar properties are placed in same column. O I II III IV V VI VII H 1 He Li Be B C N O F 4 7 9 11 12 14 16 19 Ne Na Mg Al Si P S Cl 20 23 24 27 28 31 32 35 Ar K Ca Ge As Se Br 40 39 40 75 79 80 Kr Rb Sr In Sn Sb Te I 84 85 88 115 119 122 128 127 Xe Cs Ba Tl Pb Bi 131 133 137 204 207 209 Rn (222) Used table to predict properties of undiscovered elements Eka-Silicon MM: 72 Density: 5.5g/mL Color: dirty gray Germanium MM: 72.6 Density: 5.47g/mL Color: grayish white The Modern Periodic Table Green Metals Blue Nonmetals Pink Metalloids 18 Ductile, malleable, conductive, positive ionic charge Brittle solids, often gases or liquids, negative charge Properties of both metals and nonmetals 19 Predicting Ionic Charge from the Periodic Table Alkali Metals 1A ## ## Main Group Elements Only Halogens Alkali Earth Metals Noble Gases 8A 1.0 2.0 H He 1.0 3.0 2A 4.0 3A 5.0 4A 6.0 5A 7.0 6A 8.0 7.0 9.0 4.0 10.0 Li Be B C N O F Ne 6.9 9.0 10.8 12.0 14.0 16.0 19.0 20.2 Goal: Get to column 8A by going to the right or left Right: Count each box as -1 until reaching 8A Left: Count each box as +1 until reaching 8A in previous row The correct charge is usually the smallest number Left Side (metals): Li1+ or Li-7 Mg2+ or Mg6Right Side (nonmetals): O2-, F1-, He0, Ne0 Center (metalloids): B3+ (B5-), C4+ C4- , N5+, N3Transition metals become cations: charge is unpredictable Charge is designated with a Roman numeral Iron (III) oxide contains Fe3+, Iron (II) oxide contains Fe2+ 20 Groups in the Periodic Table Alkali Metals Group 1A +1 charge Highly reactive Halogens Group 7A -1 charge Highly reactive Diatomic molecules Alkali Earth Metals Group 2A +2 charge Reactive Noble Gases Group 8A + charge (if charged) Inert Transition metals Center of table Varying (+) charge Use Roman numerals Lanthanides & Actinides Bottom of table Very reactive + charge Often radioactive 21 Elements to Memorize Know name and symbol First 4 rows of periodic table Including 1st row of transition metals Additional elements Ag: Silver Pb: Lead I: Iodine Hg: Mercury Lanthanides and actinides Know where they are in the periodic table 22 Chemical Formulas Atomic ratios in molecules Molecular Formulas Symbolic representation of molecular composition Empirical: Ratio of atoms NH2 instead of N2H4 Molecular: Types & # of atoms Structural: Relationship between atoms 23 Ionic Compounds (salts) Atoms in a reaction may gain or lose electrons Both atoms become charged and are called ions Anion: atom gaining the electron is negatively charged Cation: atom losing the electron is positively charged Will bind together to form crystals The net charge on the compound is 0 Cations & anions balance out Get charges from periodic table No distinct molecular units Positive charge of cation attracts all nearby anions 24 25 Polyatomic Ions Lose or gain electrons as a group Charge is spread over 2 or more atoms Memorize the following polyatomic ions! Ammonium Phosphate Hydroxide Cyanide Permanganate Carbonate NH4+ PO43OHCNMnO4CO32- Hydronium H3O+ Acetate CH3COONitrate NO3Sulfate SO42Chlorate ClO3Perchlorate ClO4- 26 Formula of Ionic Compounds Must match charges! 2 x +3 = +6 3 x -2 = -6 Al2O3 Al3+ 1 x +2 = +2 2 x -1 = -2 Ca2+ CaBr2 1 x +2 = +2 Br1 x -2 = -2 Na2CO3 Na+ O2- CO32- 27 Naming Compounds 28 Molecular Names Names and formulas have 2 parts, 1 for each element: Dinitrogen Tetroxide N2O4 1st word is for 1st element N: Nitrogen 2nd word is stem of element name change ending to “–ide” O: Oxygen O: Oxide Subscripts designate # of atoms, Prefixes used to name atoms: Memorize prefixes # atoms 1 2 3 4 5 prefix mono di tri tetra penta # atoms 6 7 8 9 10 N2: Dinitrogen O4: Tetroxide prefix hexa hepta octa nona deca Names of Ions and Ionic Compounds For cations: add the word ION after element name For anions: change the element name ending to –ide first Sodium Chloride: NaCl Na loses 1 eNa+ Sodium ion Cl gains 1 eNet Charge: (+1) + (-1) = 0 ClChemical Formula is NaCl Chloride ion Aluminum Oxide: Al2O3 Al (+3 charge) loses 3 eAl3+ Aluminum ion O (-2 charge) gains 2 eNet Charge: 2(+3) + 3(-2) = 0 O2Chemical Formula is Al2O3 Oxide ion 29 30 Acid Acids and Bases Arrhenius: Compound ionizes in H2O to form H+ & anions Name by changing anion –ide ending to –ic acid Add hydro to acids with HX formula: Hydrocchloric acid HCl Bronsted acids: H+ grabs H2O to form H3O+ in water Base Arrhenius: Compound ionizes in H2O to form OH- & cations Name as salts: All hydroxide salts are considered bases Bronsted base: Pulls H+ from H2O so NH3 is a base: H2O + NH3 OH- + NH4+ NH4OH Neutralization Reaction between acid & base H+ + OH- H2O and cation + anion salt HCl + NaOH H2O + NaCl (aq) Common acids and bases Memorize 31 Acids Hydrochloric Acid: Sulfuric Acid: Chloric Acid: Perchloric acid: HCl H2SO4 HClO3 HClO4 Carbonic Acid: Nitric Acid: Phosphoric Acid: Acetic Acid: H2CO3 HNO3 H3PO4 CH3COOH Bases Sodium hydroxide: NaOH Ammonium hydroxide: NH4OH Lithium hydroxide: LiOH Potassium hydroxide: KOH (actually ammonia, NH3 in H2O) Hydrates Compound is associated with a fixed number of water molecules Naming Hydrates Copper(II)sulfate CuSO4 Use numerical prefixes for # of water molecules Then add the word “hydrate” A dot shows H2O association Water molecules part of mass Copper(II)sulfate pentahydrate CuSO45H2O 32 Naming Oxoacids and their Anions: Reference Only HClO4 Perchloric acid HClO3 Chloric acid HClO2 Chlorous acid HClO Hypochlorous acid 33