section 4

Page 1

Enthalpy Change

what to know:

-enthalpy change def

-exo/endo

-breaking and making bonds

-mean bond enthalpies (accurate?)

-different types of ΔH

Enthalpy change: ΔH is the heat energy transferred in a reaction at constant pressure.

units: kJmol(-1)

ΔH° - substances are in standard states under standard conditions (1atm(100kPa), 298K)

Exo or endothermic

exo: -ΔH

-oxidation reactions e.g. CH4(g)+2O2(g)->CO2(g)+2H2O(l) ΔHc=-890kJmol(-1) and oxidation of carbohydrates, like glucose

endo: +ΔH

-thermal decomposition of CaCO3 CaCO3(s)->CaO(s)+CO2(g) ΔHc=+178kJmol(-1)

-main reactions of photosynthesis (sunlight supplying the energy

Breaking and making bonds:

-energy need to break bonds, so is endo (stronger the bond, the more energy is needed)

-energy released when bonds are formed, so is exo

ΔH for a reaction is the overall effect of these two changes

If you need more energy to break, ΔH=+, if it's less ΔH=-

Mean bond enthalpies:

-you can look them up in data books

-bond enthalpy = energy required to break bonds

-the energy needed to break a bond depends on the environment it's in, so it changes all the time

-when doing calcs, you use mean bond enthalpy - ave energy needed to break a certain type of bond over a range of compounds

ΔH of a reaction = total E absorbed - total E released

-they are AVERAGE values thf enthalpy changes aren't exact (when using mean bond enth); they are slightly less accurate than enthalpy change value calculated using Hess' Law

Different types of ΔH:

- ΔHf° : enthalpy change when 1 mole of a compound is formed from its elements in their standard states under standard conditions

- ΔHc° : enthalpy change when 1 mole of a substance is completely burned in excess oxygen to give its combustion products in their standard states under standard conditions

Page 2

Calorimetry:

what to know:

-to find out enthalpy changes

-q=mcΔT

-calorimetry in solution

-problems

Finding out enthalpy changes:

-to find out how much heat is given out during a reaction

-to find enthalpy of a combustion of a flammable liquid, you burn it inside apparatus

-as fuel burns, it heats water

-you work out heat energy by q=mcΔT

-Ideally, you want all the heat to be absorbed by water, but some will inevitably be lost to the surroundings, so it's hard to get an accurate result

Q=MCΔT

q: - heat lost/gained, j

m: mass of water, g

c: spec heat capacity of water, 4.18

ΔT: change in temp, K

NOTE: you have to divide by moles to get ΔHc and ΔHf (as they work on the basis of 1 mole)

Calorimetry in solution:

-to calculate ΔH for neutralisation, dissolution and displacement reactions

method:

1. neut: add a known vol of acid to an insulated container + measure temp

2. add known vol of alkali and record temp change

3. you can work out heat needed to raise temp using q=mc∆T (this is the heat given out by reaction)

-assume solutions have same density as water (also making it unreliable)

moles = conc x vol

Problems with Calorimetry:

all calorimetry, experimentally:

-heat absorbed by container, not water

-heat lost to surroundings (independent of insulation ability of container)

of flammable liquid calorimetry:

-combustion may be incomplete - less energy given out

-some flammable liquid may escape by evaporation (usually less volatile)

Page 3

Hess' Law

what to know:

-what it is

-to work out enthalpies of formation

-to work out enthalpies of combustion

Hess' Law: the total enthalpy change is independent of the route taken.

(to find out enthalpy changes that you can't directly find by doing an experiment, or if you're too lazy)

Formation:

(remember: ΔHof for all elements is 0)

using Hess' law: Route 1 = Route 2

products - reactants

Combustion:

reactants - products