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Honors Chemistry
Final Exam Review Guide
Chapter 6 (6-3 on):
Ionic compound ~ composed of positive and negative ions combined so +/- charges are equal
Formula unit ~ simplest collection of atoms you need to create the ionic compound's formula
Crystal lattice ~ ions in an orderly arrangement to minimize potential energy
Lattice energy ~ energy released when one mole of ionic crystalline compound is formed from the gaseous ions
Polyatomic ion ~ charged group of covalently bonded atoms
Metallic bonding ~ chemical bonding that results from attraction between metal atoms and sea of electrons (delocalized electrons that can move freely in all of the orbitals) (can be seen as communist J )
Properties of metals: malleability, absorbing various light frequencies, shininess, ductility, conducting electricity
Heat of vaporization ~ amount of heat needed to vaporize solid metal; measure of bond strength
Molecular polarity ~ uneven distribution of molecular charge, depends on polarity of each bond and molecule's geometry
VSEPR (Valence-Shell Electron-Pair Repulsion) Theory ~ states that repulsion between sets of valence level electrons surrounding an atom cause them to move as far from each other as posssible
Molecular shapes:
Linear-- 2 atoms 0 lone pairs
Bent-- 2 atoms 1 lone pair
Trigonal planar-- 3 atoms 0 lone pairs
Tetrahedral-- 4 atoms 0 lone pairs
Trigonal pyramidal-- 3 atoms 1 lone pair
Bent-- 2 atoms 2 lone pairs
Trigonal bipyramidal- 5 atoms 0 lone pairs
Octahedral 6 atoms 0 lone pairs
Hybridization ~ mixing of two or more atomic orbitals of similar energies in the same atom make new orbitals of equal energies
Hybrid orbitals ~ new orbitals made through hybridization
Intermolecular forces ~ forces of attraction between molecules
Examples: Dipole-Dipole forces: attraction between polar molecules
Hydrogen bonding: attraction of bonded hydrogen and highly electronegative atom's unshared electron pair
London dispersion forces: temporary attraction from constant motion of electrons and temporary dipoles
Chapter 7:
Monatomic ions ~ ions of a single element
If cation named by just element name
If anion named by removing ending and adding ide (ex: flouride)
Binary compounds ~ compounds made of 2 different elements
Named as name of cation, then name of anion (ex: sodium chloride)
Stock System ~ Name of cation, then roman numeral charge in parentheses, then name of anion (ex: copper (II) chloride)
Oxyanions ~ polyatomic ions w/ oxygen
Named: most common ends in -ate, one less oxygen atom than most common ends in -ite, one more than most common ends in -ate, starts in per-, two less than most common ends in -ite, starts in hypo- (ex: hypochlorite, chlorite, chlorate, perchlorate)
Molecular compounds ~ Named w/ pefixes to show numbers of atoms
Rules are: less electronegative element written first, given prefix only if more than one, followed by more electronegative element with prefix telling how many ending in -ide
Prefixes: 1 Mono- 2 Di- 3 Tri- 4 Tetra-
5 Penta- 6 Hexa- 7 Hepta- 8 Octa-
9 Nona- 10 Deca-
Covalent network compounds named just like covalent molecular compounds
Oxidation number ~ number assigned to atom in a compound to indicate the general distribution of electrons
Rules: (For covalent, but ionic can have o-numbers)
Atoms in pure element all have oxidation number 0
More electronegative element has number equal to its negative charge
Less electronegative element has number equal to its positive charge
Fluorine always has number -1
Oxygen always has -2, except in perixides (-1) or w/ halogens(+2)
Hydrogen always has +1 w/ compounds where it's less electronegative or
-1 with metals
Sum of all numbers in neutral atom must be 0
Sum of all numbers in polyatomic ion is ion's charge
Formula mass ~ sum of average atomic masses all atoms in its formula
Percentage composition ~ percent of mass of each element (not # of molecules) in a compound
Empirical formula ~ compound's formula w/ subscripts being smallest possible whole number ratios (usually same numbers as ionic formula unit, not usually right numbers in molecule)
Chapter 8:
Chemical equation ~ represents w/ symbols and formulas identities of relative amounts of reactants and products in a chemical reaction
Indications of chemical reaction:
evolution of heat and light, production of gas, formation of precipitate
Characteristics of chemical equation:
represents correct formulas for all reactants and products, law of conservation of mass satisfied
Word equation ~ where reactants and products represented w/ words + an arrow
Reversible reaction ~ reaction where products can re-form original reactants
Equations must be balanced!
Synthesis (Composition) Reaction ~ two substances combined to form new compound
Examples: metal+halogen; metal oxide+water
Decomposition Reaction ~ single compound produces two or more substances
Examples: metal oxide into metal + oxygen, metal carbonate into metal oxide + carbon dioxide, metal hydroxide into metal oxide + water, metal chlorate into metal chloride + oxygen, acids into compound + water
Single replacement reaction ~ one element replaces similar element in compound
Examples: Metal replaces metal, metal replaces hydrogen, halogen replaces halogen
Double replacement reaction ~ two compounds exchange places to make two new compounds to form gas or precipitate or water
Combustion reaction ~ substance combines w/ oxygen and releases light and heat energy (burning)
Activity series ~ list of elements organized by ease of certain reactions occurring (can replace the elements below it)
Chapter 9:
NOTE: LOTS OF STUFF IN HERE CALCULATIONS, THIS JUST GIVES DEFINITIONS!
Composition stoichiometry ~ studying mass relationships of elements in compounds
Reaction stoichiometry ~ studying mass relationships between reactants and products in a chemical reaction
Mole ratio ~ conversion factor relating amounts in moles of two substances involved in chemical reaction
Reaction under ideal conditions ~ 100% yield
Limiting reactant ~ reactant that limits amount of other reactants that combine and amounts of products that form
Excess reactant ~ substance not totally used up in reaction
Theoretical yield ~ maximum amount of product that can be produced from given amount of reactant
Actual yield ~ measured amount of product physically obtained from reaction
Percent yield ~ Actual yield x100
Theoretical yield
Chapter 10:
Kinetic Molecular Theory ~ idea that particles of matter always in motion
5 principles:
Ideal gas ~ perfect gas that fits all 5 properties above, only exists in theory
Diffusion ~ spontaneous mixing of particles of 2 substances caused by random motion
Effusion ~ process by which gas particles under pressure pass through tiny opening
Properties of gasses: expand to fill container, fluids, compressible, low density, diffuse quickly, can effuse
Real gas ~ gas that does not behave completely like KMT
WHY? B/C real gasses take up space and have attractive forces
More polar gas molecules under cooler temperatures deviate more from ideal behavior
Pressure ~ force per unit area on surface
Newton ~ SI unit for force
Barometer ~ device used to measure pressure
Pascal ~ pressure of 1 Newton per square meter
Conversion factors:
1 atm=760 mmHg=760 torr=101.3 kPa=101,300 Pa
Boyle's Law ~ volume is inversely proportional to pressure
P V=P' V'
Charles's Law ~ volume directly proportional to Kelvin temperature
V = V'
T T'
Gay-Lussac's Law ~ pressure directly proportional to volume
P = P'
T T'
Dalton's Law of Partial Pressures ~ total pressure of gas in a mixture is equal to sum of pressures of each individual gas
Chapter 11:
Gay-Lussac's law of combining volumes ~ @ constant temp and pressure, volumes of gaseous products and reactants can be expressed as ratios of small whole numbers
Avogadro's Law ~ equal volumes of gasses at same temp and pressure contain equal numbers of molecules
Standard molar volume of a gas ~ volume occupied by one mole of gas @ STP--22.4 L
Ideal Gas Law ~ mathematical relationship of pressure, volume, number of moles, and temperature of a gas PV=nRT
Constant Values: R=62.4 mmHg*L R=.0821 atm*L R=8.31 kPa*L
mol * K mol * K mol * K
Stoichiometry can be done w/ gasses plugging into the ideal gas law equation to find number of moles or molar mass
Graham's Law of Effusion ~ rate of effusion @ constant temp and pressure inversely proportional to square root of molar mass
v ² = M
v' ² M'
Chapter 12 :
Fluid ~ substance that can flow, takes shape of container
Properties of liquids: high density, incompressibility, ability to diffuse, surface tension, ability to evaporate or vaporize, ability to freeze or solidify
Crystalline solid ~ solid w/ crystals (particles arranged in orderly, geometric repeating patterns)
Amorphous solid ~ solid w/ randomly arranged particles; sometimes called supercooled liquids b/c of random arrangement of particles
Properties of solids: definite shape and volume, melting point, high density, incompressibility, extremely low rate of diffusion
Crystal structure ~ total 3-D arrangement of particles in a crystal
Unit cell ~ smallest portion of crystal lattice that shows 3-D pattern of entire lattice
Ionic crystals ~ +/- charges repeating, held together by electrostatic attraction
Covalent network crystals ~ repeating atoms covalently bonded to "nearest neighbor"
Metallic crystals ~ sea of shared valence electrons belong to crystal as whole
Covalent molecular crystals ~ covalently bonded molecules held together by intermolecular forces, strongest if molecule is polar
Equilibrium ~ dynamic condition where 2 opposing changes occur @ equal rates in a closed system
Phase ~ part of system w/ uniform composition and properties
Sublimation ~ change solid to gas
Deposition ~ change gas to solid
Le Chatelier's principle ~ when system @ equilibrium disturbed by application of stress, it attains new equilibrium to minimize stress
Equilibrium vapor pressure ~ pressure exerted by vapor in equilibrium w/ corresponding liquid @ given temperature
Volatile liquid ~ liquid that readily evaporates
Boiling ~ conversion of liquid to vapor within liquid, not just on surface
Molar heat of vaporization ~ amount of heat energy needed to vaporize one mole of liquid at boiling point
Molar heat of fusion ~ energy needed to melt one mole of solid at its melting point
Phase diagram ~ graph of pressure vs. temperature to show conditions under which each phase exists
Triple point ~ where 3 phases can coexist @ equilibrium
Critical temp ~ temp above which substance can't be liquid
Critical pressure ~ lowest pressure @ which substance can be liquid @ critical temp
Water has sp3 hybridization
Water has hexagonal particle arrangement
Water expands when it freezes (one of only compounds to do that)-- has higher density as liquid--ice floats in water
Water has fairly strong intermolecular forces (hydrogen bonding)
Water has fairly high melting and boiling points, and fairly high molar heats of fusion and vaporization