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Siddharth Nandy Period A: Mock AP Exam (Gas (Ideal Gases: imaginary gases…
Siddharth Nandy Period A: Mock AP Exam
Gas
expand to fill container
fluid
low density
compressible
Pressure: collisions of molecules with walls of container; Force/Area = Pressure
1 atm = 101325 Pa = 760 mmHg
Standard Temperature and Pressure (STP): 1 atm, 273 K
Ideal Gases: imaginary gases that perfectly fit all assumptions of the Kinetic Molecular Theory
gases are tiny particles that are far apart relative to size
only elastic collisions between particles and walls
particles in constant, rapid motion
no force of attraction between gas particles
average kinetic energy depends on temperature, not particle identity
Ideal Gas Law: PV = nRT
holes closely at P < 1 atm
P_total = n_total * RT/V
Mole fraction n_a / n_total = P_a / P_total
density = m/V = molar mass/molar volume
At STP: density = molar mass / 22.4 L/mol
D = m/V = MP/RT
Diffusion: mixing of gas
rate of diffusion = rate of gas mixing
Rate of diffusion: d_1 / d_2 = sqrt(M_2 / M_1)
Effusion: gas escaping into a vacuum
Rate of effusion: R_1 / R_2 = sqrt(M_2 / M_1)
Laws
Boyle's Law: P1V1 = P2V2
Charles's Law: V1/T1 = V2/T2
Avogadro's Law: n1/V1 = n2/V2
IMF
State of matter under given conditions determined by strength of IMFs; moderate-strong IMF = solid and liquid
Stronger IMF = higher melting/boiling points
IMF caused by opposite charges
intramolecular forces >>>> intermolecular forces (smaller distances)
Types
dipole-dipole forces
permanent polarity causes dipoles attracted to each other
London dispersion forces: temporary polarity/dipoles caused by unequal electron distribution
hydrogen bonding
highly electronegative atom leads to exposed proton that attracts electron clouds of neighboring molecules
special case of dipole-dipole force; only for atoms with H bonded to N, O, or F
ionic bonding
crystal lattice structure
metallic bonding
giant structure; electron sea model
network covalent
very strong
Factors that affect IMF
increasing size/number of electrons
fluctuations in electron distribution/density
polar molecules can induce dipoles
varies with polarizability of electrons (volume of electron cloud)
varies with shape of molecule (surface area)
ion-dipole attractions determine water solubility
vapor pressure: pressure of vapor at equilibrium
not affected by atmospheric pressure
determined by liquid's IMF strength (stronger IMF = lower vapor pressure)
Increases with temperature
Boiling point: temperature when vapor pressure = atmospheric pressure
factors into Dalton's Law of Partial Pressures
Solutions
solution: homogeneous mixture of 2+ substances
solvent: majority component; what solute is dissolved in
interact through IMFs
if UV > UU + VV, solution may not form
if UV <= UU + VV, solution will form
rule of thumb: like dissolves like
solute: minority component; what is dissolved in solvent
aqueous solution: solvent = water, solute = other solid/liquid/gas
if solution formation is endothermic but ΔH is not too large, entropy still drives solution formation
saturation: concentration
saturated: solution in which dissolved solute is in dynamic equilibrium with the solid undissolved solute
supersaturated: solution contains more than the equilibrium amount of solute
unsaturated: solution containing less than the equilibrium amount of solute
Factors that affect solubility
increasing temperature: increases (most) solids' solubility, decreases gases' solubility
increasing pressure: increases solubility of gases
Concentration measurements
Molarity: mol solute / L solution
Molality: mol solute / kg solvent
Colligative Properties: property that depends on number of particles dissolved in a solution, not particle type
vapor pressure lowering: how much vapor pressures is lowered by new solute
VPL = x_solute * P_solvent
Freezing point depression/boiling point elevation: when solute added, vapor pressure shifted downward compared to that of pure solvent --> reflects lower melting point and higher boiling point
osmotic pressure: pressure required to stop osmotic flow
increases with solute concentration
Raoult's Law: P = x * P
Modified " " : P = x_aP_a + x_bP_b
ideal solution: follows Raoult's law exactly; UV ~ UU + VV
solute dilutes solvent and ideal behavior is observed
nonideal solutions experience positive and negative deviations: vapor pressure is higher or lower than ideal
van't Hoff factor: ratio of moles of particles in solution to moles of formula unties dissolved
ideal factor only achieved in very dilute solutions
electrolyte: substance whose aqueous solution conducts an electric current
Acids/Bases
Acids: sour taste, dissolves metals, turns blue litmus paper red, neutralizes bases
carboxylic acid: R--COOH
Bases: bitter taste, slippery feel, turns red litmus paper blue, neutralizes acids
aversion to taste of bases = evolutionary adaptation again alkaloids (poisonous organic bases found in plants)
Arrhenius defintion: Acid has H+, Base has OH-
Hydronium ion: H3O+
Bronsted-Lowry definition: Acid donates proton (H+), Base accepts proton (H+)
amphoteric: substance can act as base or acid (ex. water)
conjugate acid-base pair: two substances related to each other by the transfer of proton (conjugate acid has proton, conjugate base does not)
strong acid/base: completely ionizes in solution; weak acid/base: partially ionizes in solution
strong acids: HCl, HBr, HI, HNO3, HClO4, H2SO4, HClO3
strong bases: LiOH, NaOH, KOH, Sr(OH)2, Ca(OH)2, Ba(OH)2
strength of weak acid/base determined by ionization constant; Ka=[H+][A-]/[HA], Kb = [BH+][OH-]/[B]; KaKb = Kw
use Ka and ICE table to solve for weak acid-related concentrations
5% rule: x = 0.05 * the number it is subtracted from
percent ionization: ratio of ionized acid concentration to initial acid concentration = [H+]equilibirum / [HA]initial
if dealing with two weak acids in solution, compare Ka value to see which contributes more H+
autoionization: water acts as acid/base with itself; ion product constant: Kw=[H+][OH-]=10^-14 at 25°C
acidic solution: [H+] > [OH-]
basic solution: [H+] < [OH-]
neutral: [H+] = [OH-] (= 10^-7)
pH = -log[H+]
sig. figs: X sig figs of concentration = X digits in pH mantissa
pH > 7 = basic, pH = 7 = neutral, pH < 7 = acidic
pOH = 14 - pH = -log[OH-]
pKa = -log(Ka)
Salts
conjugate of weak acid = weak base
conjugate of strong acid = pH neutral
conjugate of weak base = weak acid
Salt Solutions
Salts in which neither the cation nor the anion acts as an acid or a base form pH-neutral solutions.
Salts in which the cation does not act as an acid and the anion acts as a base form basic solutions.
Salts in which the cation acts as an acid and the anion does not act as a base form acidic solutions.
Salts in which the cation acts as an acid and the anion acts as a base form solutions in which the pH depends on the relative strengths of the acid and the base.
Polyprotic acids ionize in successive steps with different Ka values
pH usually determined by first step bc largest Ka value
Buffers
buffer resists pH change by neutralizing added acid/base
contains significant amounts of weak acid/base and its conjugate base/acid
Henderson-Hasselbach Equation: pH = pKa + log(base/acid)
equivalence point: point where moles of base = moles of acid
Ksp = solubility product constant; molar solubility = solubility in M
In general, the solubility of an ionic compound is lower in a solution containing a common ion than in pure water.
In general, the solubility of an ionic compound with a strongly basic or weakly basic anion increases with increasing acidity (decreasing pH).
Electrochemistry
Ecell = E0cell +RT/nFlog(Q) = E0cell + RT/nF * log(Q)
Anode has oxidation reaction in galvanic cell
loses mas, becomes negative
Cathode has reduction in galvanic cell
gains mass, becomes positive
electrolytic cell: cathode has reduction, is negative; anode is positive and not matches
