Physical Chemistry

Enthalpy Changes

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.

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Exothermic and Endothermic Reactions

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.
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Enthalpy Change Of...

(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.

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Hess' Law

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.

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Determining Enthalpy Changes

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.
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Bond Enthalpies

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)

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Reaction Rates

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).
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Factors Affecting Rate of Reaction (i)

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.
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Factors Affecting Rate of Reaction (ii)

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.
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The Boltzmann Distribution

(http://www.chemguide.co.uk/physical/basicrates/mbdistrib.gif) 

  • 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.
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Effect of Factors on the Boltzmann Distribution

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.

(http://www.chemguide.co.uk/physical/basicrates/mbdistrib4.gif)

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Dynamic Equilibrium

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.
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Changing the Position of the Equilibrium

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.
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Conditions

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
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The Equilibrium Constant

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.
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