b Enthalpy Changes and Thermochemistry Some important initial definitions and examples:

1.1b Enthalpy Changes and Thermochemistry Some important initial definitions and examples:

The system: The reactants and products of the reaction being studied i.e. the contents of the calorimeter.

The surroundings: The means the rest of the 'world' including the i.e. a copper calorimeter, the surrounding air etc. etc.

Enthalpy H: The heat energy content of a substance. This cannot be determined absolutely but enthalpy changes for a chemical reaction can be measured directly or indirectly from theoretical calculations using known enthalpy values.

Enthalpy change ΔH: The net heat energy transferred to a system from the surroundings or from the surroundings to a system at constant pressure. The Greek letter delta Δ in maths implies a change, in this case a net heat energy change.

ΔH = Hfinal - H_initial (the units of delta H are kJ mol^-1) or ΔH = ∑Hproducts - ∑Hreactants

or ΔH^θ(reaction) = ∑ΔH^θf(products) - ∑ΔH^θf(reactants) The Greek letter delta Δ implies 'change in' ....

The Greek letter ∑ implies 'sum of' ....'

ΔH^θf denotes a standard enthalpy of formation - which is explained further down.

Exothermic reaction

A reaction in which heat energy is given out from the system to the surroundings i.e. the enthalpy of the reacting system decreases and the temperature of the system and surroundings rises.

This means Hreactants > Hproducts so that ΔH is negative (-ve).

The enthalpy of the reaction system is decreasing.

Example: All combustion reactions are exothermic

e.g. CH4(g) + 2O2(g) ==> CO2(g) + 2H2O(l) ΔH = -890 kJmol^-1

i.e. the figure of 890 kJ released refers to the complete combustion of 1 mole of gaseous methane (24 dm³), using exactly 2 moles of gaseous oxygen (48 dm³) to form exactly 1 mole of gaseous carbon dioxide (24 dm³) and 2 moles of liquid water.

These values refer to 298K (25^oC and 1 atm/101 kPa)

Note some general points (which apply to all exothermic or endothermic changes, physical or chemical changes):

(i) All enthalpy values must be quoted with referenced to the ambient/assumed temperature and pressure of the system undergoing the physical or chemical change.

The usual standard reference conditions are 298K (25^oC and 1 atm/101 kPa), and other criteria may apply e.g. 1 molar solution if applicable.

(ii) Not only the molar quantities must clearly indicated BUT the physical states of all the substances must be clearly stated too.

This is a convenient point to make the point about the importance of state symbols via the combustion of hydrogen. eg H2(g) + ¹/2O2(g) ==> H2O(l) ΔH = -285.9 kJ mol^-1, but for

H2(g) + ¹/2O2(g) ==> H2O(g) ΔH = -241.8 kJ mol^-1

If the water forms remains as steam/vapour/gas, then 44.1 kJ less heat energy is released to the surroundings, because condensation is an exothermic process (g ==> l) and forming liquid water releases an extra 44.1 kJ. The -285.9 (~-286) kJ mol^-1 is the usual value for the enthalpy of combustion of hydrogen you will encounter in your studies because at the standard temperature of 298K water is a liquid in its normal stable state.

(iii) This sort of combustion reaction can be measured in a calorimeter (see section 1.3). BUT, however the enthalpy change is measured, all equations should be read in molar terms when dealing with enthalpy values i.e. a delta H value goes with a specific equation.

(iv) Enthalpy change values are usually quoted in kJ mol^-1, but take care in their interpretation because you must know what equation goes with the ΔH value!

eg the enthalpy of combustion usually refers to the complete combustion of one mole of the combustible material as for water above, BUT if you double the equation you must also double the enthalpy values for that equation

2H_2(g) + O_2(g) ==> 2H₂O_(l) ΔH = 2 x -285.9 = 571.8 kJ mol^-1

Endothermic reaction

A reaction in which the system takes in or absorbs heat energy from the surroundings i.e. the enthalpy of the system increases and the temperature of the system and surroundings falls OR the system must be heated to initiate the reaction and provide the heat absorbed.

This means Hproducts > Hreactants so that ΔH is positive (+ve).

The enthalpy of the reaction system is increasing.

Example: The thermal decomposition of calcium carbonate CaCO3(s) ==> CaO(s) + CO2(g) ΔH = +179 kJmol^-1

i.e. 179 kJ of heat energy must be absorbed to decompose 1 mole of solid calcium carbonate into 1 mole of solid calcium oxide and 1 mole of gaseous carbon dioxide. Mr(CaCO3) = 100, so 17.9 kJ of heat energy is absorbed in decomposing 10g of limestone. This reaction requires an experimental temperature of 800-1000^oC to achieve an appreciable rate of reaction and cannot be studied quantitatively in the laboratory. However it can be theoretically calculated from known enthalpy change values by means of a Hess's Law cycle calculation.

The two diagrams below illustrate how exothermic (left) and endothermic (right) reactions are specified on an enthalpy level diagram.

Standard conditions

Standard conditions for referencing enthalpy values are essential for communicating accurate data throughout the scientific community.

It means values measured/calculated in one laboratory/research team can be used in another scientific establishment anywhere!, OR checked for accuracy by any other scientists.

In this way accurate enthalpy data can be built up and through time validated and perhaps more accurately measured with technological developments and theoretical calculations become more reliable.

This means the reactants/products start/finish at a specified temperature, pressure and concentration whatever the 'temporary' temperature change in the reaction - which is required to calculate the enthalpy change.

The net energy change is based on the products returning to the same temperature and pressure that the reactants started at. The most frequently used standard conditions are a temperature of 298 K/25^oC (K = 273 + ^oC) and a pressure of 1 atm/101 kPa and a concentration of 1.00 mol dm^-3.

The use of standard conditions enables a database of delta H change to be assembled from which you can do theoretical calculations (see section 1.2 using Hess's Law).

Strictly speaking the standard conditions should be indicated in terms of the standard temperature and the reactants involved and standard delta H values are denoted with the Greek letter theta (θ).

By using data based on standard. agreed and defined conditions, then the data can be used universally by any laboratory around the world and also allows scientists to check each others experimental results.

Its pertinent here to consider the question - how can you have an standard enthalpy of combustion at 25^oC when the flame temperature is perhaps peaking at over 1000^oC !!!

The answer applies to all enthalpy changes what-so-ever!

The enthalpy change represents the heat energy change needed to restore the products to the temperature of the reactants at the start e.g. room temperature/25^oC.

Standard Enthalpy of Reaction ΔHr/react/reaction is the enthalpy change (heat absorbed/released, endothermic/exothermic) when molar quantities of reactants as stated in an equation react under standard conditions (i.e. 298K/25^oC, 1 atm/101kPa) Examples

(i) NaOH(aq) + HCl(aq) ==> NaCl(aq) + H2O(aq) (exothermic)

ΔH^θr,298 = -57.1 kJ mol^-1 (can also be described as an 'enthalpy of neutralisation')

(ii) CaCO3(s) ==> CaO(s) + CO2(g) (endothermic)

ΔH^θr,298 = +179 kJ mol^-1 (can also be described as an 'enthalpy of thermal decomposition')

Standard Enthalpy of Formation ΔHf/form/formation is the enthalpy change when 1 mole of compound is formed from its constituent elements with both the compound and elements in their standard states ('normal stable states) i.e. at 298K/25^oC, 1 atm/101kPa

(a) It may be endothermic or exothermic

(b) Any accompanying equation should involve the formation of 1 mole of the compound

The standard state is the most stable state at the standard temperature and pressure e.g. at 298K/25^oC and 1 atm/101kPa e.g. H2(g) H2O(l) C(s) O2(g), C3H8(g) C8H18(l) C24H50(s) CO2(g) CH3CH2OH(l) etc.

Examples

(i) C(s) + 2H2(g) ==> CH4(g) ΔH^θf,298(methane) = -74.9 kJ mol^-1 (ii) 2C(s) + 2H2(g) ==> C2H4(g) ΔH^θf,298(ethene) = +52.3 kJ mol^-1

(iii) 2C(s) + 3H2(g) + ¹/2O2(g) ==> CH3CH2OH(l) ΔH^θf,298(ethanol) = -278 kJ mol^-1 (iv) ¹/2N2(g) + O2(g) ==> NO2(g) ΔH^θf,298(nitrogen dioxide) = +33.9 kJ mol^-1 Note (a) The values can be positive/endothermic or negative/exothermic.

(b) The enthalpy of formation of elements in their standard stable states is arbitrarily assigned a value of zero.

This definition, together with experimental values of enthalpy changes allows a body of enthalpy change data to be accumulated and extended via theoretical calculations.

Standard Enthalpy of Combustion ΔHc/comb/combustion is the enthalpy change when 1 mole of a fuel (or any combustible material) is completely burned in oxygen (or air containing oxygen) equated to standard conditions (298K/25^oC, 1 atm/101kPa).

You should ensure just 1 mole of fuel appears in the equation to accompany the delta H value which is always negative i.e.

always exothermic.

Examples

(i) C3H8(g) + 5O2(g) ==> 3CO2(g) + 4H2O(l) ΔH^θc,298K(propane) = -2219 kJ mol^-1

(ii) CH3COOH(l) + 2O2(g) ==> 2CO2(g) + 2H2O(l) ΔH^θc,298K(ethanoic acid) = -876 kJ mol^-1

In the calculations explained below just the subscripted letters r/f/c will be used for brevity and a temperature of 298K and a constant pressure 1atm assumed unless otherwise stated. There is more the enthalpies of combustion of alkanes and alcohols in section 1.4a

Standard enthalpy of neutralisation is the energy released when unit molar quantities of acids and alkalis completely neutralise each other at 298K (pressure effects are insignificant for reactions only involving liquids/solutions/solids) (i) NaOH(aq) + HCl(aq) ==> NaCl(aq) + H2O(l) ΔH^θneutralisation = -57.1 kJ mol^-1

(ii) Ba(OH)2(aq) + 2HNO3(aq) ==> Ba(NO3)2(aq) + 2H2O(l) ΔH^θneutralisation = -116.4 kJ mol^-1 (iii) ¹/2Ba(OH)2(aq) + HNO3(aq) ==> ¹/2Ba(NO3)2(aq) + H2O(l) ΔH^θneutralisation = -58.2 kJ mol^-1

Note! It looks as if the enthalpy of neutralisation of barium hydroxide is approximately double that of sodium hydroxide ie ~ twice as exothermic! Well yes it is! and no it isn't!

Yes - ~twice as much energy is released per mole of soluble base/alkali.

No - however, on the basis of heat released per mole of water formed, they are actually very similar.

In other words, which value you quote, depends on which point you want to make.

Yet another example of carefully qualifying enthalpy values with respect to the context.

More on enthalpies of neutralisation Bond Enthalpy ('bond energy')

This is the average energy absorbed to break 1 mole of a specified bond when all species involved are in the gaseous state.

e.g. for (i) H2(g) ==> 2H(g) ΔH = +436 kJ mol^-1 for the H-H bond

or for (ii) CH3CH2Br(g) ==> CH3CH2(g) + Br(g) ΔH = +276 kJ mol^-1 for the C-Br bond

It is always endothermic and the reverse process - bond formation, is always exothermic. In many cases the values are averaged from a variety of 'molecular' situations. More on this in the bond enthalpy section.

Some examples of points made on this page with reference to an enthalpy level change diagram

General points: Arrows pointing downwards represent exothermic changes and arrows pointing upwards represent endothermic changes