Chemistry EOY

What are the ideal gas assumptions?

constant random motion

combined volume of all gas molecules is negligible compared to container volume

negligible IMF of attraction and repulsion forces between gas molecules

collisions between gas molecules are perfectly elastic (average KE constant at constant T)

manometer

open tube: Pgas = Patm + pgh

closed tube: Pgas = pgh

collecting gas over water

must first level water level inside and outside

Ptotal = Patm = Pgas + Pwater vapour

Deviation from ideal gas (graph)

positive deviation (PV/nRT >1): high pressure, significant volume

negative deviation (PV/nRT <1): low temperature, significant IMF

Ideal gas law and molar mass (equations, charles’ law, boyle’s law, avogrado’s law)

PV = nRT

M = ρRT/P

charles’ law: V = kT

boyle’s law: V = k/P

avogrado’s law: V = kn

Dalton’s law of partial pressure (equations)

Pa = Xa,gas * Ptotal

Ptotal = Pa + Pb + Pc

Kinetic molecular theory (maxwell boltzmann curve, KE, v_rms)

Maxwell boltzmann curve: number or frequency of molecules against average speed or energy. area under graph constant

High temperature (high KE) = wider range lower peak, peak shift right (higher average speed)

Low temperature (low KE) = smaller range, higher peak, peak shift left (lower average speed)

KE = (mv^2)/2

v = sqrt(3RT/M) - average KE affected by molar mass and temperature

Diffusion (what is the definition, what is it affected by)

gradual mixing of molecules of gas 1 with molecules of gas 2

higher speed (high temperature, low molecular mass) = higher rate of diffusion

Effusion (what is the definition, what is it affected by)

lower molecular mass = effuse more rapidly

move more quickly, higher chance to hit hole

Vaporisation

reverse of condensation (same ΔH magnitude)

molecules escape liquid and enter gaseous state

energy required to overcome IMF, endothermic (ΔH>0)

Condensation

reverse of vaporisation (same ΔH magnitude)

high pressure, low temperature = IMF > KE = liquid

exothermic (ΔH<0)

Evaporation

some molecules at the surface with high KE overcome IMF of attraction to break free from neighbouring molecules and enter gaseous phase

increased rate at higher temperature

maxwell boltzmann curve shift right and wider = more particles have enough temperature to overcome IMF

open container vaporisation

evaporate to dryness

vapour diffuse from liquid surface

lower tendency to condense back, rate of evaporation > rate of condensation

closed container vaporsation

equilibrium state, no net change

rate of evaporation = rate of condensation

evaporated gas molecules hit surface and reenter liquid

vapour pressure

pressure exerted by vapour when liquid and vapour in dynamic equilibrium

high IMF, low temperature = low vapour pressure

Boiling point

vapour pressure = external pressure

bubbles can form

higher altitude = lower external pressure = lower boiling point

Phase diagram

solid-liquid interface: melting point at different pressures

positive slope as higher pressure = lower volume = higher density = need more energy to break free from 3D solid (except water)

liquid-vapour interface: boiling point at different pressures

triple point: all three states coexist in equilibrium

critical point: above critical point, liquid and gas phase indistinguishable

energy used when heating substance

constant temperature: energy used to overcome IMF of attraction to change state

increasing temperature: energy used to increase average KE of molecules

Viscosity (definition, affected by)

resistance to flow, ease of movement with which molecules move relative to each other

large size, great strength/extensiveness of IMF = higher viscosity

lower temperature = less molecules with sufficient KE to overcome IMF = higher viscosity

more viscous means harder to flow

surface tension (definition, explanation)

only for surface molecules

tendency of liquid to maintain minimum surface area, or energy required to increase surface area by unit amount (J/m^2)

molecules at surface only interact with molecules beneath = net inward attractive force on surface molecules = surface molecules more tightly packed

cohesive and adhesive forces (concave and convex)

cohesive: IMF with similar molecules within liquid

adhesive: IMF with different substance

concave = adhesive > cohesive

convex = cohesive > adhesive

capillary action (definition, explanation)

rise of liquids up very narrow tubes

low cohesive, high adhesive = sucked up tube (higher adhesive = greater capillary action)

net upward attraction until gravity balances out forces

solution

homogeneous mixture of 2 or more substances in a single phase (can be solid, gas, aqueous)

solvent = larger amount

solute = smaller amount

Molarity

moles of solute per Litres of solution

changes with temperature

molality

moles of solute per kg of solvent

more stable

mole fraction

moles of solute per moles of solution

mass percentage

mass of solute over mass of solution x 100%

parts per million

mas of solute over mass of solution x 10^6

Solubility

grams of solute per 100g solvent

favourable if energy released from solvent-solute bond forming greater than energy required to overcome pure solute and pure solvent IMF

Solubility of ionic compounds

soluble in water: energy released by strong IDF overcome ionic bonds

form hydrated ions

dissolved solute in dynamic equilibrium with solid

insoluble in non-polar: energy by weak IDF insufficient

solubility rules: all nitrates, sulfates (except Pb, Ag, Ca, Ba), chloride (except Ag, Pb)

carbonates, oxides, hydroxides not soluble (except group 1, ammonium)

solubility of covalent compounds

insoluble in water: LDF and mainly HB greater than DDF

soluble in non-polar: comparable weak LDF = solvent penetrated solute

Liquids

miscible = homogeneous 1 phase solution

like dissolves like - similar/stronger IMF when mixed

if molecule is polar and non-polar, as non-polar part gets longer, solubility decreases as polar part becomes smaller and molecule behaves more like non-polar, with stronger LDF

Saturated solution (explanation)

maximum solute in solution in dynamic equilibrium with solid

generally higher temperature = higher solubility

Gas solubility

higher molar mass - larger electron cloud = stronger DIDF = greater solubility

higher temperature = difficult to form IMF with solvent = lower solubility

Henry’s law

solubility of gas in liquid is directly proportional to gas pressure

Solubility (M) = kP (k = Henry’s law constant, different for each gas)

higher pressure = greater frequency of gas molecules colliding with liquid surface = higher chances of reentering solution = increased solubility

S1/S2 = P1/P2

limitation: assumes gas does not interact with solvent

explaining gas solubility scenarios

canned drinks: packed under high pressure - more CO2 gas soluble

can opened = pressure decrease = solubility decrease = CO2 escape

underwater: high pressure = more gases soluble in blood

rapid ascent = pressure decrease = solubility decrease = gas form bubbles in blood

bubbles below boiling point: other gases dissolved in liquid that become less soluble at higher temperatures

Effect of increased non-volatile solute concentration on Vapour pressure of solvent

vapour pressure decrease as less solvent molecules at surface, so number of molecules that can escape at given time decreases while number of molecules that can redissolve remains unchanged

note: if ionic solvent dissolves into 2 molecules, effect is doubled

Raoult’s law

Pa = Xa,liquid Pºa

Ideal solution vapour pressure in solution of 2 or more volatile liquids

Ptotal is straight line between Pºb and Pºa

assume no change in IMF when solute is added

linear following raoult’s law

real close to ideal when solute concentration low or solute and solvent similar IMF and size

Real solution vapour pressure positive deviation

solute-solvent interactions weaker than pure solute and pure solvent

vapour pressure higher than predicted as more molecules escae readily

Real solution vapour pressure negative deviation

negative deviation: solute-solvent interactions stronger than pure solute and solvent

vapour pressure lower than expected

Boiling point elevation by adding solute

lower vapour pressure = higher temperature required to raise vapour pressure to 1 atm

ΔTb = iKm (final boiling point = normal boiling point + Tb)

Freezing point depression by adding solute

solute disrupt IMF between solvent molecules - solvent must lose more energy - freezing point decrease

ΔTf = iKm (final freezing point = normal freezing point - Tf)

Van hoff factor i

number of particles added per unit of solute

not all dissociated at any point in time (form ion pairs when ions of opposite charge collide)

i approaches maximum at lower concentration (ions further apart)

higher charge = stronger attraction = lower actual i

Factors affecting frequency of effective collisions

Surface area

particles per unit volume (concentration or partial pressure)

temperature - increased frequency and increased energy

catalyst - pathway with lower Ea (lower energy transition state)

rate law

rate = k[A]^m[B]^n

units for k change depending on order of reaction

overall order = m + n

Zero order reaction

rate = k (horizontal line)

[A] = -kt + [A]o (linear)

First order reaction

rate = k[A] (linear)

[A] against t is inverse graph (curve)

half life constant = ln2/k

ln[A] = -kt + ln[A]o (when a = 1)

Second order reaction

rate = k[A]^2 (quadratic graph)

[A] against t is curve down (similar to first order)

1/[A] = kt + 1/[A]o (positive gradient)

Pseudo-order

all reactants kept in large excess/constant except for one

change in concentration of others negligible

rate = k[A][B] = k’[A] (k’ = k[B])

rate constant k

k = Ae^(-Ea/RT), R = 8.314J/Kmol

lnk = (-Ea/RT) + lnA

Elementary steps

sum up to form overall chemical reaction

rate law of elementary step follows stoichometry ratio

overall rate law follows rate law of slow step, same k

catalyst = reacted and reproduced, intermediate = not present as reactant, produced and used up

unimolecular = 1 reactant molecule, bimolecular = 2 reactant molecules

Homogeneous catalysts

same phase as reactants

lower energy transition state

Heterogeneous catalyst

different phase from reactants

limited sites, become saturated = zero order

Dynamic equilibrium

rate of forward reaction = rate of backward reaction

reversible reaction in closed system

Equilibrium constant

Kc = [products]/[reactants]

no units, remember to raise to power of stoichometry, not reaction mechanism

only affected by temperature, not initial amount

concentration or partial pressures at equilibrium

pure solid/liquid not included

Kp = Kc(RT)^(n) where n = (c+d)-(a+b)

Position of equilibrium

Kc > 1 - position of equilibrium to the right (more products than reactants)

Kc < 1 - position of equilibrium to the left (more reactants than products)

Reaction quotient Q

Q = K - equilibrium, no net reaction

Q < K - [product] lower, [reactant] greater than at equilibrium, net reaction in forward direction

Q > K - [product] greater, [reactant] lower than at equilibrium, net reaction in backward direction

Le Chatelier’s principle

when stress applied, system adjusts position of equilibrium to counteract change

no change in equilibrium position if catalyst added

if more water added - all concentrations decrease equally - try to increase number of particles (compare if more reactants or more products)

adding inert gas at constant volume has no effect

Explaining effect of change of reversible system

state change and effect on Q

use LCP to describe how system reacts (direction favoured and why)

shift in equilibrium position

Graph

sharp change observed for change in concentration, partial pressure and volume, followed by gradual change until equilibrium

no sharp change for temperature, only gradual change (K changes)

Bronsted Lowry acids and bases

acid: proton donor, forms conjugate base

base: proton acceptor, forms conjugate acid

conjugate acid-base pair differ by only 1 H+

Ka and Kb

Ionisation of acid or base in water

Ka x Kb of conjugate base = Kw = 10^14

Higher Ka or Kb means stronger acid or base

pKa = -lg Ka (same for Kb)

pH and pOH

pH = -lg[H+]

pOH = -lg[OH-]

pH + pOH = pKw = 14

Strong acids and bases

ionise completely in H2O, strong electrolyte

use ICF table to find pH

very weak conjugate acid/base

strong acids: Hcl, H2SO4, HNO3, HBr, HI, HClO3, HClO4

strong bases: oxides and hydroxides of group 1 and 2 metals

Weak acids and bases

ionise partially in H2O, weak electrolyte

use ICE table and Ka or Kb to find pH

significant conjugate acid/base

start reacting after strong acid/base finishes reacting

percentage ionisation