section 4

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

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=MCΔTq: - heat lost/gained, jm: mass of water, gc: spec heat capacity of water, 4.18ΔT: change in temp, KNOTE: 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 reactionsmethod:1. neut: add a known vol of acid to an insulated container + measure temp2. add known vol of alkali and record temp change3. 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)

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 2products - reactants

Combustion:reactants - products