AP Chemistry Flashcards Flashcards

Terms Definitions
(1)
Conversions between Temperatures
°C à K add 273 (no negative temperatures; absolute scale)
K à°C subtract 273 (negative temperatures are possible)
Kelvin must be used with the gas laws to avoid negative
volumes, pressures, and moles (impossible) and also to provide the correct proportional relationships between temperature and a gas's pressure and volume (K.E. directly proportional to absolute temperature)
Temperatures used in thermochemistry (heats of reaction and
calorimetry) are ∆T therefore it makes no difference which scale is used since one degree Celsius is the same size unit as one Kelvin.
(2)
Heat Capacity and Specific Heat
Heat capacity: the amount of energy required to raise the temperature of an amount of a substance by 1°C; an extensive physical property
(dependent on sample size)
Heat capacity has the units J°C
Specific Heat: the amount of energy required to raise the temperature of 1 gram of a substance by 1°C; an intensive physical property
(independent of sample size)
Specific heat has the units J g°C
Heat capacity or specific heat will be negative if defined as amount of energy lost to reduce temperature by 1°C. The greater the specific heat of a substance, the slower the substance heats up or cools down.
(3)
Electromagnetic Radiation
Electron Transitions and Photons Emitted
Kelvin.
c = λν
(c = 3.00 x 108 m s-1)
(h = 6.63 x 10-34 J s)
(A = 2.18 x 10-18 J)
(n = main energy level)
(where nb > na)
∆E can be converted to frequency or wavelength using one or
both of the first two equations above.
(4)
Solution by Dilution
When preparing diluted solutions from more concentrated stock
solutions (solution by dilution), the moles of solute are equal in
both solutions, therefore:
M1V1 (concentrated) = M2V2 (diluted)
Where all volumes represented are final volumes. To find the
amount of water needed to prepare the diluted solution, simply
subtract the final volume of the concentrated solution needed
from the final volume of the diluted solution made (V2 - V1)
(5)
Quantum Numbers
Principal (n): main energy level; values range from 1 à ∞
Azimuthal (l): number of subshells and orbital shape; values range
from 0 àn-1; 0 = s (spherical), 1 = p (dumbbell), 2= d, 3 = f
Magnetic (ml): number of spatial orientations of orbital
Spin (ms): axial rotation of electron; values are either + ½
or - ½
If n=1, then l and ml are automatically 0, typical set: 4, 1, -1, + ½ means electron is in 4th main energy level, in a p-orbital, in one of its 3 possible orientations (probably px), spinning clockwise.
(6)
Rules for Electron Configuration
Aufbau principle: an electron will occupy the lowest possible
energy orbital that will receive it; orbitals overlap in their
energies at higher main energy levels and should be filled
according to the diagonal rule or by reading the periodic table.
Hund's rule: when filling orbitals of equal energy (including
hybrids) electrons of the same spin will occupy each orbital as
singles before pairing or when filling a subshell with more than
one orbital, place one electron in each orbital before pairing.
Pauli exclusion principle: no two electrons of the same atom
will have the same exact four quantum numbers (spins will be
opposite)
(7)
Monatomic Ions
Group I, II, III metals: 1+, 2+, 3+ respectively
Group V, VI, VII nonmetals: 3-, 2-, 1-, respectively, when forming
binary ionic compounds; -ide ending when naming these anions
Transition metals with only one charge: Ag+, Zn2+, Cd2+,, Bi3+
Transition metals with more than one charge: Cu+/Cu2+, Fe2+/Fe3+,
Co2+/Co3+, Mn2+/Mn3+, Sn2+/Sn4+, Pb2+/Pb4+
Stock system of naming: using Roman numeral in () for cation's charge
Old system: Latin root +-ous for lower charge /+ -ic for higher
charge (i.e. cuprous, cupric; ferrous, ferric; stannous, stannic, etc)
(8)
Polyatomic Ions
NH+ (ammonium); H3O+ (hydronium); C2H3O2- (acetate);
NO3- (nitrate); MnO4- (permanganate); ClO4- (perchlorate);
ClO3-, BrO3-, IO3- (chlorate, bromate, iodate); CN- (cyanide)
OH- (hydroxide); HCO3-, HSO4- (bicarbonate, bisulfate)
CO32- (carbonate); SO42- (sulfate); C2O42- (oxalate);
CrO42- (chromate); Cr2O72- (dichromate); O22- (peroxide);
PO43- (phosphate); one less oxygen: -ate ion becomes -ite ion
(9)
Drawing Lewis Structures
1. Count total number of valence electrons from all atoms
(Positive polyatomics, subtract charge; negative polyatomics,
add charge)
2. Draw skeletal structure and connect all ligands to central
atom with single bond. Subtract electrons used for this from
the total.
3. Place remaining electrons in pairs starting with the ligands
and moving to central atom until octets are complete.
4. Check for octets. If atoms lack octets, form multiple bonds,
one at a time, only until all octets are complete.
Exceptions to octet rule: H, Be, B (less than 8) or Period 3 and higher nonmetals (can be greater than 8 because of d-subshell)
(10)
Bond Order and Bond Properties
Bond order is number of bonds between two (bond order of 1 for single, 2 for double, etc); for resonant structures, the B.O. will almost always be fractional and can be determined by the total bonds divided by the number of area where the bonds appear.
Bond length (or distance) is distance between the nuclei of two atoms covalently bonded when the P.E. is at the minimum; indirectly related to bond order.
Bond energy is the energy required (+kJ/mol) to break one mole of covalent bonds; the opposite of bond energy (-kJ/mol) is the energy given off when one mole of those bonds are formed; directly related to bond order as is vibrational frequency which represents frequency of bending and oscillating motions of bonds.
(11)
VSEPR Theory and Molecular Geometry
MX2: linear (180°)
MX3: triangular planar (120°)
MX2E: bent, V-shaped (<120°)
MX4: tetrahedral (109.5°)
MX3E: trigonal pyramidal (107°)
MX2E2: bent (105°)
MX5: trigonal bipyramidal (90°, 120°)
MX4E: distorted tetrahedral
MX3E2: T-shaped (90°)
MX2E3: linear (180°)
MX6: octahedral (90°)
MX5E: square pyramidal
MX4E2: square planar (90°)
(12)
Valence Bond Theory
1. Covalent bonds are formed by orbital overlap.
2. The greater the overlap, the stronger the covalent bond.
3. When bonding, atoms position themselves such that
maximum orbital overlap can occur (hybridization)
MX2: sp-hybrid
MX3, MX2E: sp2-hydrid
MX4, MX3E, MX2E2: sp3-hybrid
MX5, MX4E, MX3E2, MX2E3: sp3d -hybrid
MX6, MX5E, MX4E2: sp3d2- hybrid
(14)
Acidic and Basic Anhydrides
Acidic anhydrides: nonmetallic oxides (gases) that react with water to form acidic solutions: oxidation states remain unchanged.
Examples: CO2, SO2, SO3, N2O3, N2O5, P4O6, P4O10
SO3(g) + H2O(l) à H2SO4(aq)
Basic anhydride: metallic oxides (solids) that react with water to form basic (alkaline) solutions.
Examples: Group I and Ca, Br, Sr metallic oxides are soluble and therefore produce the most basic solutions.
Na2O(s) + H2O(l) à 2NaOH(aq)
(15)
Arrhenius Acids and Bases
Br∅nsted-Lowry Acids and Bases
Lewis Acids and Bases
Arrhenius acid: any substance that increases hydronium ion
concentration in aqueous solution.
Arrehenius base: a soluble hydroxide that increases hydroxide ion concentration in aqueous solution.
Br∅nsted-Lowry Acid: a substance that donates protons to a Br∅nsted- Lowry base.
Br∅nsted-Lowry Base: a substance that accepts protons from a Br∅nsted-Lowry acid
(Br∅nsted-Lowry acid-base reactions have conjugate acid-base pairs.)
Lewis acid: substance that accepts e-pair in forming coordinate bond.
Lewis base: substance that donates e-pair to form coordinate bond. (Lewis acid-base reactions yield coordination cmpds/complex ions)
(16)
Periodic Trends in Strength of Acids and Bases
The strength of acids is determined by how easily the hydrogens on the molecule are ionized by water to form hydronium ions.
Binary acids: stronger down a group due to larger atoms; stronger across a period due to increasing electronegativity of nonmetal.
Oxoacids: stronger across a period due to increasing electro negativity of central atom (if number of lone O's the same); stronger with more lone O's (if central atom is same) due to lone O's pulling electron density away from H's on the molecule make them easier to ionize.
Soluble ionic oxides or hydroxides (Group I, Ca, Sr, Ba) are the
strongest bases; less basic across periods toward amphoterism in
middle and more molecular in nature; increasing acidity towards end of
period.
(17)
Assigning Oxidation Numbers
1.Oxidation state of a free element is zero
2.Oxidation state of Group I, II, III metals: +1, +2. +3,
respectively
3.Oxidation state of F is always -1
4.Oxidation state of H is +1 except in metallic hydrides where
it&#039;s -1
5.Oxidation state of O is -2 except with F where it&#039;s +2, in
peroxides where it&#039;s -1, or in superoxides where it&#039;s -1/2
6.Oxidation state of a monatomic ion is equal to ion&#039;s charge
7.Sum of oxidation states of all atoms in a polyatomic ion is
equal to the ion&#039;s charge.
8.Sum of oxidation states of all atoms or ions in formula of a
compound equals zero
(18)
Balancing Redox Equations
1.Assign oxidation numbers to each atom in the equation.
2.Identify the substance oxidized (oxidation state increases) and the substance reduced (oxidation state decreases) and divide into half- reactions (oxidation half-reaction and reduction half-reaction)
3.Balance identical atoms in each half-reaction
4.Balance charge on each side of each half-reaction by adding
electrons to appropriate side (added e's will be opposite for each half)
5.Apply appropriate factor to each half-reaction in order to equalize number of electrons. Factor out electrons
6.Combine half-reactions to get net reaction (equation)
7.Check that number of atoms of each element and total charge on the left side equals number of atoms and total charge on right.
(19)
Redox Reactions in Acidic and Basic Environments
In acidic and basic environments:
1. Separate reaction into two half-reaction (oxidation and
reduction)
2. Balance all atoms except O's and H's
3. Add one water for each deficit O to the deficit side
4. Add one hydrogen proton for each deficit H to the deficit side
5. Balance charge in each half-reaction by adding electrons
6. Equalize number of electrons in both half-reactions with
factors
7. Recombine half-reactions to form balanced net equation
Continue with the following for basic (alkaline) environment:
8. Add one hydroxide to both sides for each hydrogen proton
9. Combine hydrogen protons and hydroxides to form water
10. Simplify by factoring out like particles on both sides of net
(20)
Solubility Rules
1.All Group I metallic salts, ammonium salts, nitrates, acetates, chlorates, and perchlorates are soluble.
2.All chlorides, bromides, and iodides are soluble except when
with silver, lead(II) or mercury(I)
3.All sulfates are soluble except when with lead(II), Ca, Sr, or
Ba.
4.Group I, Ca, Sr, Ba metallic oxides are soluble (basic
anhydrides) as well as their resulting hydroxides
5.Insoluble (except with cations in Rule 1): carbonates,
phosphates, sulfides and bisulfites (hydrogen sulfites)
(21)
Common Laboratory Oxidizing Agents
Oxidizing agents are reduced in a redox reaction:
MnO4-(aq) à Mn2+(aq)
(acidic)
MnO4-(aq) à MnO2 (s)
(basic)
Cr2O72-(aq) à Cr3+(aq)
(acidic)
CrO42-(aq) à Cr(OH)3 (s)
(basic)
Also, all strong nonmetals are excellent oxidizing agents:
O2 (g) à O2- (combustion if fast, corrosion if slow)
Cl2 (g) à Cl-- (free halogens reduce to halide anions)
(22)
Born-Haber Cycles
(∆Hf for Ionic Compounds)
1st ionization energy
+ electron affinity (-)
bond energy (+)
heat of formation
lattice energy (-) + 12
The standard heat of formation for the compound NaCl (s) is the sum of all energies
involved and the net result must be negative in order for the compound to be stable. Since
all energies are in kJ mol, be aware of when 1 2 mole of reactant is used. Also, add 2nd
ionization energy for Group II metal reactants.
(23)
Common Laboratory Reducing Agents
Reducing agents are oxidized in a redox reaction:
HSO3-(aq) à SO42-
(acidic)
SO32-(aq) à SO42-
(basic)
S2O32-(aq) à SO42-
(acidic w/ Cl2)
S2O32-(aq) à S4O62-
(acidic w/ I2)
Fe2+(aq) or Sn2+(aq) à Fe3+(aq) or Sn4+(aq)
Also, all strong metals are excellent reducing agents:
Na (s) à Na+
Ba (s) à Ba2+ (free metals oxidize to metallic cations)
(24)
∆Hf from Bond Energies
1.Start with a balanced equation showing the Lewis structure of the reactants and products.
2.Determine the moles of reactant bonds broken (+ bond
energies) or moles of final gaseous atoms (+ atomization
energies). Multiply by number of moles as factors.
3.Determine the moles of product bonds formed (- bond
energies). Multiply by number of moles as factors.
4.The net energy change (∆Hf) will be the sum of all energies
from #2 and #3 above.
(25)
Solution Concentrations
Molarity (M): moles of solute per liters of final solution
g solute
M = mol solute
= molar mass
L solution
L solution
Molality (m): moles of solute per kilogram of solvent
g solute
m = mol solute = molar mass
kg solvent
kg solvent
mole fraction (X): moles solvent (or solute)/ total moles of both
X = mol solvent (or solute)
Total mol of both
(26)
Percent Composition
Empirical and Molecular Formula Determination
Percent composition is by mass and can be determined from a formula or from the mass data from an analysis. It is the ratio of masses (part over whole) times 100 and is constant for any size sample of a pure substance.
The empirical formula of a compound can be determined from either mass percent data or from mass data. From the percent data, think of percentages as grams/100g cmpd. Convert g to mol and use answers as initial subscripts in formula. Find smallest whole-number ratio by dividing all subscripts by the smallest.
The molecular formula is determined by how many times the empirical molar mass goes into molecular molar mass and then increasing empirical formula by any finding >1.
(27)
Hess's Law of Heat Summation
If a series of equations are additive such that they yield a net equation then the heats of reaction associated with each equation are also additive and yield the net heat of reaction for the net equation.
When manipulating thermochemical equations:
1. If the equation is reversed, the sign of the heat is made
opposite.
2. If the equation is increased/decreased by a mathematical
factor, the same factor is also applied to the heat.
Hess's law of heat summation (if heats of formation are given):
∆H° = Σ∆Hf° (products) - Σ∆Hf° (reactants)
(∆Hf free element = 0)
(28)
Chemical and Volumetric Analysis
Chemical analysis involves finding the mass percent of a
particular substance in a compound or in a mixture of pure
substances. Volumetric analysis involves a titration. One
needs to know the mass of the original sample to be analyzed
and the balanced equations for all reactions involved in the
analysis procedures. Using coefficients from balanced equations (reaction stoichiometry) and subscripts from chemical formulas (composition stoichiometry), one should be able to determine the mass of the substance being sought. The mass percent will be this mass divided by the mass of the original sample times100. This percent data or the mass data can then be used to determine the empirical formula for an unknown compound.
(29)
∆Hf from Combustion Data
To find ∆Hf° of a substance using its combustion data (∆Hc° and the combustion equation):
1. Write balanced combustion equation of the substance so
moles of carbon dioxide and water are known (combustion
products).
2. ∆Hf° of the substance equals the sum of the ∆Hf° of all
combustion products minus ∆Hc° of the substance (given).
∆Hf° = Σ∆Hf° (products of combustion) - ∆Hc°
(30)
Metathesis Reactions
1. Precipitate: know solubility rules to determine the
formation of a precipitate in a metathesis reaction.
2. Gas: Look for formation of obvious gases such
as H2S or acidic anhydrides in weak acids such
as CO2- in H2CO3, SO2 in H2SO3, N2O3 in HNO2;
also NH3 in the weak base NH4OH.
3. Weak electrolyte: look for all of #2 above as
well as water (one product of an acid-base
neutralization) or HC2H3O2.
(31)
History of Periodic Table Development
Cannizarro: made possible experimental atomic masses by reactions
Newlands: when elements listed by increasing atomic mass, every 7th element had similar properties
Mendeleev: arranged elements by increasing atomic mass such that elements with similar properties fell in same row. Blanks left for undiscovered elements and properties predicted according to periodic law he first stated. Noble gases not yet discovered (unreactive).
Ramsay: discovered noble gases thus adding new group of elements
Moseley: using x-ray deflection angles, he discovered number of protons in atoms (atomic number). New periodic law developed based on arrangement of elements by increasing atomic number.
(32)
Periodic Trends in Atomic Properties
Atomic radii decrease within period as additional protons increase nuclear charge without additional shielding. Increases within group as electrons occupy higher energy levels further from nucleus and additional shielding by inner-shell electrons. Lanthanide contraction results in atoms in Period 6 being the same size as Period 5 because of large increase in protons (f-subshell) without increase in effective
shielding.
Ionization energy increases within period from greater nuclear charge holding electrons more tightly. Decreases within group because electrons more loosely held. Successive I. E.'s always increase with greatest occurring at noble gas configuration. Cation smaller than atom.
Electron affinity has greatest negative values (most stable anion) at halogens. Trends hard to identify. Anion larger than neutral atom.
(33)
Dalton's Atomic Theory
Discovery of Subatomic Particles
Dalton's atomic theory postulates that matter is made up of tiny,indivisible particles called atoms, which can't be created nor destroyed. Atoms of the same element are identical; atoms of different elements are different. Atoms combined in small whole-number ratios to form compounds. In reactions, atoms are combined, separated, rearranged.
Thomson: discovered electrons using cathode-ray tube. Concluded they are negative and have a small relative mass ratio and that there must be positive particles to balance charge of electrons and make up most of mass of atoms.
Millikan's experiment determined charge of electrons
by amount of current to keep oil drops (with e 's) suspended.
Rutherford: discovered atom's nucleus with alpha-particles and gold foil. Concluded most of the atom is empty space and the nucleus is very small.
(34)
Fundamental and Derived Quantities
Only seven fundamental quantities are used in measuring:
1. length (l) measured in SI units of meters (m)
2. mass (m) measured in SI units of kilograms (kg)
3. time (t) measured in SI units of seconds (s)
4. temperature (T) measured in SI units of kelvins (K)
5. amount of a substance measured in SI units of moles (mol)
6. luminous intensity measured in SI units of candelas (cd)
7. electrical current measured in SI units of amperes (A)
Other quantities used in measuring are mathematical combinations of the above and are therefore derived quantities. Examples: velocity is 1/ t, volume is l3, area is l2, density is m/l3, joule is m l2/t2, etc.
(35)
Unit Conversions and Factor-Label
Unit conversions are made using factor-label method (dimensional analysis) and knowing meanings of metric prefixes:
Mega (M) = 106; kilo (k) = 103; deci (d) = 10-1; centi (c) = 10-2;
Milli (m) = 10-3; micro (µ) = 10-6; nanao (n) = 10-9 ; pico(p) = 10-12
Conversions are made by using an equality (1 m = 100 cm) as a ratio (1m/100 cm or 100 cm/1m) called a conversion factor. If multiple quantities involved (mi/hr to m/s), convert each quantity separately.When quantities involving exponents are converted (in3 to cm3), use regular conversion factor (2.54 cm/ 1 in) and then cube the entire conversion including the numbers.
(36)
Density and Specific Gravity
Density : an intensive physical property of pure substances defined as mass per unit volume. Volume of regularly-shaped solid is found by application of a formula to its dimensions. Volume of irregularly-shaped solid is found by submerging completely in water and finding out what volume is displaced (1 mL = 1 cm3)
Specific gravity: a unitless ratio of the density of a pure substance to the density of water in the same units. If this ratio is known, then the density of a substance can be found in any unit as long as a table of the density of water in those same units is available.
(37)
Significant Figures
The number of significant digits in a measurement indicates the precision of the instrument making the measurement. All nonzero digits are significant.
Rules governing whether a zero is significant are:
1. Zeros between significant digits are significant.
2. Zeros after a nonzero number and after a decimal are significant.
3. Zeros preceding a written decimal are significant.
4. Zeros preceding an understood decimal (not written) or the first nonzero number are not significant.
Rules for rounding for operations involving significant figures:
1. Addition/subtraction: round to least number of decimal places.
2. Multiplication/division: round to least number of significant
3. Log x: number of sig. figs. in x = number of decimal places in log x.
(38)
Naming Compounds (including Acids)
Ionic compounds: name of compound is the names of ions together with cation first. Roman numeral in parentheses is part of cation name if &gt;1 oxidation state. Name of monatomic anion ends in -ide. When writing formulas from name, write skeletal formula first and balance charges by adjusting subscripts (&gt;1 polyatomic ion in parentheses)
Binary covalent: name of compound is the names of elements with - ide ending for 2nd element (most electronegative). Prefix used with 1st element if &gt;1 atom; prefix used with 2nd element for any number of atoms. Hydrogen compounds named as ionic unless in aqueous solution (acids).
Binary acids: hydro- + root of nonmetal + -ic acid (aq)
Oxoacids: per -ate ion = per -ic acid; -ate ion = -ic acid; -ite ion = -ous
acid; hypo - ite ion = hypo- ous acid (no hydro -prefix with oxoacids).
(39)
Limiting and Excess Reactant
Percent Yield
If > 1 reactant is given in stoichiometry problem, one limits amounts of products formed. To determine limiting reactant, convert to moles,divide by coefficient of each from balanced equation, and least will be limiting. The original given amount of limiting is then used to determine the amount of products formed (theoretical yield).
To determine leftover reactant, subtract the amount of excess reactant used to form the product from its original given amount. For molarity of leftover, divide leftover moles by additive volumes (in L) of solution mixture.
Percent Yield is actual yield divided by theoretical yield times 100.
Percent yield is usually < 100% (never >100%) due to loss of product (careless technique or damaged equipment) or side reactions which produce undesirable products (contaminated, dirty equipment)
(40)
Writing Net Equations
1.Solid reactants and products (ppt) will always appear in net
2.Gaseous reactants and products will appear as molecules in net
3.Free elements (or oxides and water) as reactants will produce a single compound as the product. Be sure to write the correct formulas.
4.A single reactant will decompose into simpler elements or
compounds (not ions). Be sure to write the correct formulas.
5.Two aqueous ionic reactants will be either metathesis or redox (if acidic/basic environment specified). If metathesis, look for ppt, gas, or weak electrolyte. If redox, work with net ions only.
6.Strong acids and soluble compounds appear as ions in the net.
7.Free element and aq ionic as reactants will be single-replacement
8.In acid/base neutralization, if salt is soluble, net is H+ + OH-
=H2O
(41)
Ideal Gas Equation
PV = nRT
where R = 0.0821 L atm mol-1 K-1 = 62.4 L torr mol-1 K-1
(molar mass may be found this way for molecular formula
determination from empirical formula) by molar volume (22.41)] [gas density at STP is molar mass divided
(42)
Ion concentrations and pH / pOH
Kw = [H3O+][OH-] = 1.00 x 10-14 M2
(H3O+ same as H+)
In dilute aqueous solutions of strong acids and bases, if the
concentration of one of the ions is known, then the concentration of the
other can be found using the above relationship.
The pH scale goes from values of 0 - 14 with 7 being neutral. Values of 0 - &lt;7 are acidic and values &gt;7 - 14 are basic (alkaline). (pOH scale is the opposite.) pH + pOH = 14 for strong acid/ base solutions.
pH = -log[H3O+]
pOH = -log[OH-]
[H3O+] = antilog (-pH)
[OH-] = antilog (-pOH)
(43)
Graham's Law of Effusion
Graham's Law states that the ratio of rates of effusion of two gases is the square root of the inverse of their molar masses. The ratio of rates may also be expressed as ratio of velocities or distance traveled. The molar masses ratio may also be expressed as densities ratio.
Rate of effusion of gas A/Rate of effusion of gas B = Square root of (MM gas B/ MM gas A)
Molar mass of an unknown gas determined in this way can be used in determination of a molecular formula from an empirical formula.
(44)
Ka for Weak Acids
Kb for Weak Bases
Weak acid (HA) at equilibrium: HA ßà H+ + A-
Ka = [H][A]/[HA]
Weak base (B) at equilibrium: B + H2O ßà BH+ + OH-
Kb = [BH][OH]/[B]
For a weak acid and its conjugate base (or weak base and its conjugate acid): Kw = Ka x Kb = 1.00 x 10-14 M2
(45)
Kinetic Molecular Theory
Postulates of the kinetic molecular theory of matter:
1. All matter is made up of particles that are in constant motion
2. Particles of a gas move at high speeds in random straight
lines and collisions between particles and/or container walls
are elastic (energy may be transferred but net energy remains
constant).
3. A gas's pressure results from collisions from its container
walls.
4. The average kinetic energy of particles is directly
proportional to absolute (K) temperature.
This theory is useful in explaining the properties and behavior of gases, liquids and solids.
(46)
Hydrolysis of Salts
In salt formed from strong acid / strong base, neither ion hyrolyzes therefore the pH of its solution is 7.
In salt formed from weak acid / strong base, the cation hydrolyzes: BH+ + H2O ßà B + H3O+ , and its solution is acidic (pH<7). Use Ka.
In salt formed from strong acid / weak base, the anion hydrolyzes: A- + H2O ßà HA + OH- , and its solution is basic (pH>7). Use Kb.
In salt formed from weak acid / weak base both ions hydrolyze and pH of its solution will be acidic if Ka > Kb or basic if Kb>Ka.
(47)
Binding Forces and Properties of Crystals
Ionic crystals: cations and anions at lattice sites; electrostatic forces of attraction between oppositely-charged particles; hard, brittle, high relative m.p./b.p., conductors only in molten state or (aq); NaCl etc.
Metallic crystals: metallic cations at lattice sites surrounded by "sea of electrons" (mobile electrons); electrostatic attractions; range of hardness, m.p./b.p. (dependent on charge and size), malleable, ductile, good conductors of electricity as solids; Na, Al, Cu, etc.
Covalent Molecular: Atoms/molecules at lattice sites; van der Waals intermolecular attractions; soft, low relative m.p./b.p.; nonconductors (except for covalent acids); H2O, CO2, H2S, C6H12O6, etc.
Covalent network: atoms/molecules at lattice sites; covalent bonds between all particles; hard/brittle, very high relative m.p./b.p.; SiO2
(48)
Relationship between Kc and Kp
Kp = Kc (RT)∆ng
Where ∆ng = moles of gaseous products - moles of gaseous reactants and R = 0.0821 L atm mol-1 K-1
(49)
van der Waals Intermolecular Forces
London dispersion forces: Weaker; exist between all molecules and caused by induced dipoles on the molecule from movement of electrons; stronger in more complex molecules (higher MM or longer C-chains) because more electrons moving and more dipole sites.
Dipole-Dipole attractions: Stronger; exist between polar molecules
only; the greater the dipole moment, the stronger the attraction.
H-bonding: A particularly strong dipole-dipole between H of one
molecule and either F, O, or N on another molecule.
Relative melting/boiling points and heats of fusion/vaporization are higher in substances with stronger van der Waals forces.
(50)
pH of Buffered Solutions
Buffered acid: Made by combining a weak acid and a salt containing the acid's conjugate base anion (i.e. HNO2 and NaNO2)
Buffered base: Made by combining a weak base and a salt containing the base's conjugate acid cation (i.e. NH3 and NH4Cl)
(61)
Acid-Base Titrations
Strong Acid/ Strong Base : pH determined by excess [H+] or [OH-]; pH at equivalence point is 7, because neither ion of the salt hydrolyzes.
Strong Acid/Weak Base : changing pH during titration determined by buffer equation and changing ratio of unionized base (decreases by added H+) to conjugate acid cation (increases by salt); pH at equivalence point is greater than 7 because of hydrolysis of the salt cation; use cation's Ka from base Kb .
Weak Acid/Strong Base : changing pH during titration determined by buffer equation and changing ratio of unionized acid (decreases by added OH-) to conjugate base ion (increases by salt); pH at equivalence point is < 7 because of hydrolysis of salt anion; use anion's Kb from acid Ka.
(62)
Standard Cell Potential (E°cell)
E°cell = E°reduced - E°oxidized
If E°cell is positive, cell reactions are spontaneous as
written.
If E°cell is negative, cell reactions are spontaneous in
reverse.
E°- are standard reduction potentials of the substances
involved.
(63)
Solubility-Product Constant (Ksp)
Ksp represents the equilibrium state of a saturated solution. To
determine Ksp, write the dissociation equation for the salt and make sure solubility data is in mol L (may have a convert g, 100mL, H2O, etc)
Given Ksp value, solving for "x" from the Ksp --expression gives the molar solubility of the salt.
(64)
Relationship between Kc and E°cell
E­cell = 0.0592v/n log Kc = 0.0257v/n ln Kc
log Kc­ nE/0.0592v or ln Kc­ = nE/0.0257v
(65)
Common Ion Effect and Solubility
The common ion effect decreases the molar solubility of a salt by increasing molar concentration of an ion common to the salt thus shifting solution equilibrium to the left (according to LeChatelier).
AgCl(s) ßà Ag+(aq) + Cl-(aq)
Kap = [Ag+][Cl-]
[Cl-] will be the concentration of the common
ion initially present (as in NaCl).
Solving for [Ag+] as "x" will give the new molar solubility of the
slightly soluble salt with the common ion initially present in the
solution.
(66)
Relationship with ∆G° and E°cell
Delta G = -n(Faraday)E of cell
where n = moles of e- in the balanced redox equation and = faraday (F) = 96,500 C/ mol e-
(67)
Complex Ion Formation and Solubility
The formation of a complex ion increases the molar solubility of a salt because the ion involved is removed from solution thereby shifting solution equilibrium to the right (according to LeChatelier).
AgCl(s) = Ag+(aq) + Cl-(aq)
Ksp
Ag+(aq) + 2NH3(aq) = Ag(NH3)2+(aq)
Kform
AgCl(s) + 2NH3(aq) ßà Ag(NH3)2+(aq) + Cl-
Kc {add]
(solve for "x" to find new molar solubility of salt)
(68)
Nernst Equation
The Nernst equation relates changes in E for a cell to concentrations of cell solutions (or partial pressures) if not standard (1 M, 1 atm) or equal :
E­cell = E­cell - log Q = E­cell - ln Q
Predict changes in E by LeChatelier before calculating to verify answer.
(69)
Electrolytic Cells
Galvanic Cells
In both cells, anode is where oxidation takes place and cathode is where reduction takes place. Charges are opposite between each cell.
Electrolytic cell : cathode is negative because voltage source supplies electrons here so metallic cations are reduced to atoms; anode positive therefore attracts electrons away from metallic atoms thereby oxidizing them to cations. With electrolysis of solutions, remember competing oxidation-reduction half-reactions for water.
Galvanic (Voltaic) cell : cathode is positive because metallic cations with highest reduction potential take electrons from here, leaving an "electron deficit"; anode is negative because metallic atoms (lowest reduction potential) deposit electrons here, leaving it "electron rich". Flow of electrons in external circuit is always from anode to cathode.
(70)
Types of Radiation
Alpha rays (): nucleus; low penetrability and low energy; minimal damage
Beta radiation ():-; high speed electrons (neutron to proton); more energy, higher penetrability; x-rays
Gamma rays (): high energy EM radiation (no particle); greatest energy and penetrability
Neutron: 0
Proton: +
Positron: +
(51)
Raoult's Law
Raoult's Law predicts the lowering of vapor pressure of a solvent upon the addition of a solute. Solute particles at the surface mean less solvent particles can evaporate, therefore reducing vapor pressure.
Nonvolatile solute: Psoln = ( Xsolv)(P°solv)
X = mole fraction of solvent; P° = vapor pressure of pure solvent
Volatile solute: PT = Xa P°a + Xb P°b
X = mole fraction of part; P° = vapor pressure of pure part
Positive deviation: solute-solvent attractions weaker than each pure part
Negative deviation: solute-solvent attractions stronger than each pure
part
(52)
Collision Theory
Collision Theory is used to explain changes in rates of reactions (not
shifts in equilibrium) and postulates the following :
1) In order for reactions to occur, particles must be in proximity
to each other such that collisions can occur
2) Collisions between reactant particles must be effective in
order to result in product particles. Effective collisions
require that the reactant particles have proper orientation
to each other and that they collide with enough energy
(activation energy) to result in bond-breaking in reactants and
bond-forming in products.
This theory is used to explain why concentration of reactants, surface
area, temperature, and catalysts affect the rate of a reaction.
(53)
Colligative Properties
The addition of a solid solute to a solvent lowers the freezing
point and raises the boiling point of the solvent. This is a
colligative property because it depends solely on the number of
particles rather than the nature of the particles.
Nondissociating nonelectrolytes:
∆ tf = (molality) (Kf)
f.p. = "normal" f.p. - ∆tf
∆ tb = (molality) (Kb)
b.p. = "normal" b.p. + ∆tb
(Kf for H2O is 1.86° C molal: Kb for H2O is 0.512° C molal)
Electrolytes: the colligative effect is 2, 3, or 4 times that of above assuming 100% ionization because electrolytes dissociate into more moles of ions (particles) than nonelectrolytes (find in the formula)
(54)
Arrhenius Equation
The Arrhenius Equation is used to determine activation energy (Ee- ) or changes in rate (k) with changes in absolute temperature (T):
ln(k1/k2)=Ea/R(1/T2-1/T1) or ln(k2/k1)=Ea/R(T2-T1/T1T2)
When R = 8.314 J mol-1 K-1
(55)
Gibb's Free Energy and Spontaneity
Gibbs free energy equation relates spontaneity of reactions to their ∆H°, ∆S°, and Kelvin temperature. In order for a reaction to appear spontaneous, ∆G° must be negative in: ∆G° = ∆G° - T∆S°
When ∆H° is negative and ∆S° is positive, both natural tendencies are satisfied therefore reaction always appears spontaneous.
When ∆H° is positive and ∆S° negative, neither natural tendency is satisfied therefore reaction never appears spontaneous.
When ∆H° and ∆S° are the same sign, spontaneity is temperature - dependent. If both positive, appears spontaneous at high T; if both negative, appears spontaneous at low T. To find the T at which the reaction appears spontaneous, set ∆G° = 0 and solve for T.
(56)
Relationship ∆G° and Kc /Kp
∆G = ∆G° + RT (lnQ) where R = 8.314 J mol-1 K-1
Q is the mass action quotient found at any point in the reaction.
If ∆G is negative, the reaction will continue in the forward
direction to reach equilibrium; if ∆G is positive, the reaction
will continue in the reverse direction to reach equilibrium.
∆G° = - RT (ln Kp) (for reactions involving gases)
∆G° = - RT (ln Kc) (for reactions involving solutions)
The K values found in this way are called thermodynamic equilibrium
constants and you are restricted to finding Kc or Kp (or using them)
depending on whether the reaction involves gases or solutions.
(57)
Changes in Concentration over Time
Half Life
Changes in concentrations of reactants over time (kinetics):
ln([A]o/[A]t)=kt or ln([A]t/[A]o)= -kt
If the concentration of reactants decreases by 50%, then the
ratio of concentrations equals 1/0.5 = 2 and ln2 = 0.693. The
time required to decrease concentration by ½ is called the half-
life of the reaction. It can be found by rearranging the above
and solving for "t". This is the same half life used in nuclear
chemistry and measures decay:
t1/2 = 0.693/k where k is the rate or decay constant
(58)
The Gas Laws
Boyle's law : at constant T, gas V is inversely proportional to P.
PV = k
P1V1 = P2V2
Charles' law : at constant P, gas V is directly proportional to T.
V/T=K V1/T1=V2/T2
Gay - Lussac's law: at constant V, gas P is directly
proportional to T.
P/T = K P1/T1=P2/T2 P1T2=P2T1
Ideal gas equation : relates gas P, V, T and n (moles)
PV = nRT
(59)
Standard Heat of Reaction
Standard Entropy Change
Standard Free Energy Change
∆H° = Σ ∆H°f (products) - Σ ∆H°f (reactants)
(∆H°f of a free element is 0)
∆S° = Σ S° (products) - Σ S° (reactants)
(S° is never 0 for a substance except at 0 K)
∆G° = Σ ∆G°f (products) - Σ ∆G°f (reactants)
(∆Gf° of free element is 0)
(60)
Dalton's Law of Partial Pressures
Dalton's Law of Partial Pressure : At constant volume, the total
pressure of a mixture of gases is the sum of the partial pressures
of each gas in the mixture.: PT = P1 + P2 + P3 + P4 +...
If the volume or temperature changes when mixing the gases,
apply Boyle's or Gay-Lussac's Law to get the new partial
pressure of each gas at new volume or new temperature and then
add to get total pressure.
Dalton's Law is used when a gas is collected by water
displacement in order to remove water vapor pressure (at the
collection temperature) to get pressure of the "dry" gas. Total
pressure is atmospheric pressure as long as water levels are
equalized.
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