Physical Chemistry

Chemical energy

Is a special form of potential energy that lies within chemical bonds - forces of attraction that bind atoms together in a compound.

Enthalpy changes

• During a reaction, bonds break in reactants and form in products.
• This changes the chemical energy.
• Bonds break = exothermic (-)
• Bonds form = endothermic (+)

Enthalpy

Heat content stored in a chemical system.

Enthalpy changes

Is the head energy transferred in a reaction.

Law of conservation of energy

"If heat is released, the amount of energy that leaves a chemical system is exactly the same as the amount that goes to the surroundings. No heat energy is lost."

Exothermic reactions

• Enthalpy of products is smaller than the ethalpy of reactants.
• Heat loss to surroundings.
• Negative enthalpy change.

Endothermic reactions

• Enthalpy of products is greater than the enthalpy of reactants.
• Heat gain to chemical system.
• Positive enthalpy change.

(standard conditions: 100kPa, 298K)

Standard enthalpy change of reaction ΔrHo

Enthalpy change when the reaction occurs in the molar quantities shown in the equation.

Standard enthalpy change of formation ΔfHo

Enthalpy change when 1 mole of a compound is formed directly from its elements.

Standard enthalpy of combustion ΔcHo

Enthalpy change when 1 mle of a substance is completely burned in oxygen.

Standard enthalpy change of neutralisation ΔneutHo

Enthalpy change when an acid and alkali react together.

Activation energy is the minimum energy required for a reaction to take place.

Hess' Law

Total enthalpy change is independent of the route taken.

• Enthalpy of Formation
  • F = products - reactants
• Enthalpy of Combustion
  • C = reactants - products
• It is not always possible to measure the enthalpy change directly: high Ea, slow reaction and more than one reaction can take place.
• Breaking bonds = endothermic as energy is needed.
• Forming bonds = exothermic as energy is released.

Bond dissociation enthalpy is the amount of energy needed per mole (g)

Average bond enthalpy is the energy needed to break 1 mole of bonds in gaseous states over many different compounds.

Determining enthalpy changes

Q= mCΔT

• Q = heat lost/gained.
• m = mass of water (g)
• C = specific heat capacity - 4.18
• ΔT = change in temperature.

There may be different values with experimental data and text books...

• Incomplete combustion
• Heat loss to surroundings
  • To get a more accurate result, use a bomb calorimeter.

Bond enthalpies

Enthalpy change that takes place when breaking 1 mole of a given bond in the gasous states.

• Endothermic
  • Energy needed to break bond : bonds formed stronger than bonds broken.

• Exothermic
  • Energy needed to form bond : bonds broken are stroner than bonds formed.

ΔH = Σ(bond enthalpy of bonds broken) - Σ(bond enthalpy of bond formed)

Rate of reaction

The change in the concentration of products and reactants in given time.

Rate = change in concentration    moldm-3 s-1

                      time

Collision theory

States that for a reaction to take place molecules must:

• Collide with the minimum amount of kinetic energy.
• Collide in the right direction (orientation).

1) Concentration

• Increase in conc = increase in ROR.
• More molecules in the same vol - more frequent and successful collisions.

2) Pressure

• Increase in pressure of gaseous reactants = increase in conc = increase in ROR.
• Same number of molecules occupy a smaller volume - more frequent and successful collisions.

3) Temperature

• Increase in temp = increase in ROR.
• More kinetic energy = more molecules overcoming activation energy - collide more successfully.

4) Surface Area

• Using smaller particles increases SA = increase in ROR.
• Allows more frequent and successful collisions.

5) Catalyst

Lowers the activation energy of a reaction by providing an alternative route. e.g. transition metals - Fe in haber proccess or enzymes.

• More particles overcome Ea.
• Less energy, less fossil fuels burnt, less CO2 released.

• Area under the curve = the number of molecules in the sample.
• No molecules in the system with 0 energy.
• No maximum energy for a molecule.
• Only molecules with energy under the Ea are able to react.

1) Temperature

• Kinetic energy of molecules increases.
• Area under the curve remains the same.
  • Flattens and shifts to the right.

2) Catalyst

• Ea is lowered.
  • More molecules in the system overcome the Ea.

Dynamic equilibrium

The rate of forwards and backwards reaction are the same - so the conc of products and reactants are the same. (reached in reversible reactions)

Le Chatelier's Principle

When a system in dynamic equilibrium changes, the equilibrium will shift to oppose the change.

The position of the equilibrium can be changed by...

• Concentration of reactants/products.
• Pressure.
• Temperature.

1) Concentration

• Increase in reactants = shifts to the right. (RR)
  • To decrease conc of reactants.
• Increase in products = shifts to the left.
  • To decrease conc of products.

2) Pressure

• Increase in pressure = shifts to the side with fewer moles.
  • In order to decrease the pressure.

3) Temperature

• Increase in temp = adding heat = moves to the endothermic side.
  • To absorb heat.
• Decrease in temp = removing head = moves to the exothermic side.
  • To replaces heat.

• Catalysts help reach the equilibrium faster - NOT changing the position.

Conditions

• You have to compromise...
  • A low temp = slower rate.
  • A high pressure = too expensive and dangerous.

Haber Process

• Temp = 400-500°C.
• Pressure = 200 atm.
• Catalyst = Iron.

Making ethanol

• Temp = 300°C.
• Pressure = 60-70 atm.
• Catalyst = H3PO4

The Equilibrium Law

Tells us the relative proportions of reactants and products present in an equilibrium.

aA + bB <--> cC + dD        [ABCD] = concentration. [abcd] = balancing numbers.

Kc = ([D]^d) x ([C]^c)

       ([A]^a) x ([B]^b)

• If Kc is 1 = halfway between the reactants and products.
• If Kc is smaller than 1 = products favoured.
• If Kc is bigger than 1 = reactants favoured.

• Forward reaction = endothermic.
  • As temp increases, Kc increases.
  • Conc of products increase, conc of reactants decrease.
• Backward reaction = exothermic.
  • As temperature increases, Kc decreases.
  • Conc of products and reactants decrease.