HL Chem Semester 2

Exothermic

Surroundings get hotter, negative value, makes bonds
ex. combustion

Endothermic

Surroundings get colder, positive value, breaks bonds

Bond Enthalpy

Enthalpy change when 1 mole of covalent bonds in a gaseous molecule is broken under standard conditions

∆H Atomization

1 mole of gaseous atom is formed from an element. Always endothermic. (s) → (g)

IE

Electron is removed from an isolated atom in the gaseous state. Always endothermic

EA

an electron is added to an isolated atom. 1st EA is exothermic, second is endothermic

∆H latt (Lattice Enthalpy)

1 mole of an ionic compound is broken apart into its constituent gaseous ions. Always endothermic
Depends on:

• The higher the charge, the higher the mp, the higher the enthalpy

• The smaller the ionic radius, the higher the enthalpy

High charge + low radius = higher charge density

∆H sol

1 mole of solvate is dissolved in excess solvent. Endothermic or Exothermic

∆H hyd

1 mole of gaseous ion is surrounded by water molecules to form a solution. Always exothermic. Same rules as ∆H latt.

Factors that affect reaction rates

1. the nature of the reactant: varies based on reactivity

2. concentration: higher concentration, more collisions, higher rate

3. Temperature: higher temp, more kinetic energy, more collisions, higher rate

4. Catalysts: changes mechanism and lowers activation energy

5. Surface area: a solid in a powder for increases in surface area which increases the rate

Measuring the rate of reaction

• reactions that produce gas: collect gas and measure its volume/pressure

• reactions involving ions: conductivity can be measured (probe, voltage)

• reactions that change colour: intensity of colour can be measured with a spectrophotometer

Equilibrium: Concentration change

equilibrium shifts in the opposite direction of a concentration increase

Equilibrium: Pressure

Increase pressure → equilibrium shifts to decrease pressure

Q (Reaction Quotient) vs Kc

Q can take any concentration, Kc is only equilibrium concentrations

Rule of 100

divide initial concentration by Kc. >100 assumption to ignore 0 can be made.

Arrhenius Theory

Acids produce H+ when dissolved in water, bases produce OH- when dissolved in water

Brontsed-Lowry Theory

Acids are proton (H+) donors, bases are proton acceptors

Amphoteric Substances

Can act as acids or bases ((H2), HCO3- (baking soda))

Amphiprotic

Can accept or donate a H+ (H2O)

Lewis Theory

Acid is an electron acceptor, base is an electron donour

A conjugate base is

negatively charged

A conjugate acid is

positively charged

Strong acid

reaction goes to completion, conjugate base is weak. Strong Ka, will only dissociate slightly
Ex: HCl, HBr, H2SO4, HNO3, HPO4, HCLO4

Weak Acid

established an equilibrium, conjugate base is strong
Ex: carboxylic acids,HF, H2CO3, H2S, H3BO3

Monoprotic acids

only contain 1 ionizable H (important for ionization energy)

Strong bases

Weak conjugate acids. ex: group 1 or group 2 + OH

Weak Bases

Strong conjugate acids. ex: anything with nitrogen

Salt hydrolysis

when salt reacts with water. does not occur in salts that form neutral solutions.

The larger the Ka

the lower the pka, the stronger the acid

More positive the entropy

the higher the spontaneity, the higher the disorder
increases with temperature and gaseous state molecules

Moving a system to greater stability by:

decreasing enthalpy and increasing entropy

At equilibrium (∆G)

∆G = 0, but G does not = 0

Spontenous at high temps

∆G = + / - , ∆H = +, ∆S = +

Spontaneous at low temps

∆G = + / - , ∆H = -, ∆S = -