CCEA ADVANCED SUBSIDIARY

CHEMISTRY

 

MODULE 1
 

1.6    Energetics

 

The concept of an enthalpy change, ΔH, for endothermic and exothermic changes. Simple enthalpy level diagrams. Explanation of enthalpy changes associated with changes of state.

Thermochemistry

 

The branch of chemistry concerned with the energy changes that occur during chemical reactions and phase changes is called thermochemistry.  Practically all chemical reactions are accompanied by enthalpy (heat) changes.  This means that the total energy of the products (H2) is different from the total energy of the reactants (H1).  The enthalpy change is given by

 

ΔH = H2 – H1       sometimes called the heat of reaction

 

Exothermic and endothermic reactions

Exothermic

During the chemical reaction, heat energy is released to the surroundings. The energy loss from the chemicals provides the energy gain for the surroundings and so the temperature increases. The enthalpy of the products is lower than the enthalpy of the reactants  (H1 > H2) and ΔH is negative. The reaction is exothermic.

 

 

     heat

                                                                                          H1           reactants

 


                                                                               enthalpy                               ΔH (negative)

heat                  hotter than                       heat                                         

                         surroundings                                         H2                                   products

 

 


                                                                                                         reaction path

    heat

 

 

Endothermic

During the chemical reaction, heat energy is taken in from the surroundings. The energy gain for the chemicals is provided by the energy loss of the surroundings i.e. the temperature decreases. The enthalpy of the products is greater than the enthalpy of the reactants (H1 < H2) and ΔH is positive.  The reaction is endothermic.

                                  heat

                                                                                          H2                                    products

                                                                                                                                 ΔH (positive)

                          colder than

heat                 surroundings                      heat           H1              reactants

 

 


                                                                                                         reaction path

 

                                  heat

 

Enthalpy is measured in joules (J) and kilojoules (kJ).

 

Exercise 1

Classify the following changes as exothermic or endothermic

(i)                 Sodium hydroxide dissolves in water and the temperature of the solution rises.

(ii)               Ammonium chloride dissolves in water and the temperature of the surroundings drops.

(iii)             Hydrogen and oxygen combine explosively to form water.

(iv)              Liquid water condenses to ice at 0oC.

(v)                Liquid nitrogen (boiling point = 77K) boils spontaneously at room temperature.

 

 

Enthalpies of phase changes

When a chemical changes from one state to another (a phase change) energy is either released or has to be supplied.

                   fusion                               vapourisation


SOLID                                 LIQUID                              GAS   


                                  sublimation

 


The enthalpy of fusion ΔHfusion is the heat absorbed when 1 mole of a solid is converted to a liquid at the melting point. This energy is needed to overcome the attractive forces between the particles in the solid e.g. in ice energy is needed to break down the H-bonded structure of the solid in order to form liquid water.

ΔHvapourisation is a measure of the attractive forces between particles in a liquid. Energy must be supplied to separate the particles in a liquid in order to form the gas phase.

The energy required to form 1 mole of a gas starting from the solid is called the enthalply of sublimation ΔHsublimation. As can be seen from the diagram this is simply the sum of the enthalpies of fusion and vapourisation.

 

ΔHsublimation  = ΔHfusion  + ΔHvapourisation

 

 

Simple experimental study of enthalpy changes and calculation of standard enthalpy changes from experimental data.

Experimental determination of enthalpy changes

In discussing enthalpy changes we have spoken of heat transfers to and from the surroundings and of being able to measure them. In practice, it is much easier to measure the enthalpy change of a reaction using a calorimeter in which the reaction is insulated from the surroundings. Then, for an exothermic reaction, the energy which would have been given to the surroundings is trapped resulting in an increase in temperature of the chemical system.

If the maximum temperature change of the system is measured and recorded and if the heat capacity of the system is known it is possible to calculate the enthalpy change for the reaction.

There are several types of insulated calorimeter which can be used for simple experiments as shown below. These in effect prevent heat change with the surroundings so that all energy changes occur within the system. The energy exchanged with the surroundings is usually small enough to be ignored.


 


Before you begin experiments the example below will show how you can calculate the enthalpy change of a reaction from experimental data.

 

Worked Example

An excess of zinc powder was added to 50.0 cm3 of 0.100 M AgNO3 in a polystyrene cup.  Initially, the temperature was 21.10 o C and it rose to 25.40 oC.

Calculate the enthalpy change for the reaction

 


Zn(s) + 2Ag+(aq)                              Zn2+ (aq) + 2Ag(s)

 


Assume that the density of the solution is 1.00 g cm-3 and its specific heat capacity is 4.18 kJ kg-1 K-1.  Ignore the heat capacity of the metals.

 

Solution

 

1. Since the polystyrene cup is an insulator and its heat capacity is almost zero, you can assume that no energy is exchanged between system and the surroundings. All the chemical energy released in the reaction is transformed into heat energy which raises the temperature of the solution. The total energy change in the system is zero, so you can write..

 

[Enthalpy change, ΔH,      +          [change in heat        = 0

  due to reaction ]                           energy of solution]                          

 

Since      

[change in heat        = mass x specific heat capacity x temperature change =mcp ΔT

  energy of solution]

 

you can now write: ΔH + mcp ΔT = 0

 

∴ ΔH  = -mcp ΔT =    -50.0 kg  x  4.18 kJ kg-1 K-1  x  4.30 K  =  -0.899 kJ

                                  1000

 

2.  The value -0.899 kJ is the enthalpy change for the amounts used in the experiment. To obtain a value for the enthalpy change of reaction, compare the amounts used in the experiment with the amounts shown in the           equation:

 


Zn(s) + 2Ag+(aq)                              Zn2+ (aq) + 2Ag(s)

                             1 mol     2 mol

The amount of silver ion used = 0.0500 dm3 x 0.100 mol dm-3  = 5.00 x 10-3 mol    

 

∴ the enthalpy change using 2 mol of Ag+                                             

=  -0.899 kJ x    2.00 mol             =  -360 KJ                                 

    5.00 x 10-3 mol

 

 

3.  Now write the complete thermochemical equation:


Zn(s) + 2Ag+(aq)                                    Zn2+ (aq) + 2Ag(s) : ΔH = -360 kJ mol-1

 

 


Note that enthalpy changes related to equations, which include all standard enthalpy changes, have the unit kJ mol-1.

 

See experiment sheets for;

Determination of the enthalpy of a reaction

 

Standard enthalpy changes of reaction, combustion, formation, neutralisation.

Standard Enthalpy Changes

The enthalpy changes that occur during a reaction vary depending on the temperature, pressure, the physical state of the substance and the amount of the substance involved.

The standard molar enthalpy of a reaction is the enthalpy change per mole of the reaction at 298K (25oC) and 1 atm. pressure.  ΔHθ (298K)

e.g.


(a) 2H2 (g) + O2 (g)               2H2O (l)      ΔHθ(298K) = -571.6 kJ mole-1


(b) H2 (g) + ½ O2 (g)                          H2O (l)      ΔHθ (298K) = -285.8 kJ mole-1


(c) H2 (g) + ½ O2 (g)               H2O (g)     ΔHθ (298K)= -241.8 kJ mole-1

 


The difference between (b) and (c) is the standard enthalpy of vaporisation of water.

 

Enthalpy changes are given the general name enthalpy of reaction but are usually classified according to the type of reaction.

 

(a)  Standard Enthalpy of Formation ΔHθf (298K),

This is the enthalpy change when a mole of substance is formed from its elements in their standard state.

 


C(s) + 02 (g)                            C02 (g)                        ΔHθf  = -393.5 kJ mole-1


½ H2 (g)  +  ½ Cl2 (g)                HCl (g)            ΔHθf  = -92.3 kJ mole-1


2Al(s) + 3/2 02(g)                 Al203(s)          ΔHθf  = -1676 kJ mole-1

 


The more negative the value the more stable the compound is relative to its elements.

 

 

 

(b)  Standard Enthalpy of Combustion ΔHθc (298K),

This is the enthalpy change when one mole of a substance in its standard state is completely burned in air or oxygen at 298K.

e. g.


C3H6 (g) + 9/2 O2 (g)         3CO2 (g) + 3H2O (l)                       ΔHθc = -2291.7 kJ mole-1


CH4 (g) + 2O2 (g)                CO2 (g) + 2H2O (l)              ΔHθc = -890.2 kJ mole-1

 


Enthalpies of combustion are always negative as combustion is an energy releasing (i.e. exothermic) process.

 

(c)  Standard Enthalpy of Neutralisation ΔHθn (298K),

This is the enthalpy change per mole of water formed in the neutralisation of an acid by an alkali.

e.g.


HCl (aq) + NaOH (aq)                   NaCl (aq)  +  H2O (l)                   ΔHθn = -5 7. 1 kJ mole-1


HNO3 (aq) + KOH (aq)                    KNO3 (aq)  +  H2O (l)        ΔHθn = -5 7. 1 kJ mole-1


½ H2SO4 (aq) + NaOH (aq)             ½ Na2SO4 (aq)  +  H2O (l) ΔHθn = -5 7. 1 kJ mole-1

 


If spectator ions are removed it can be seen that the term enthalpy of neutralisation refers to the reaction :


           H+ (aq)  +  OH- (aq)                           H2O (l)                 ΔHθn = -5 7. 1 kJ mole-1

We might therefore expect that the neutralisation of any acid by any alkali would give the same value for the enthalpy of neutralisation. This is certainly true for the reaction between strong acids and alkalis, shown above, which are completely dissociated into ions. However weak acids and alkalis are only partially ionised in water and some energy is required to complete their dissociation before reaction between hydrogen ions and hydroxide ions can occur. This means that less energy is available as heat and the enthalpies of neutralisation are less than the value quoted above.


e.g. CH3COOH  +  NaOH                       CH3COONa  +  H2O  ΔHθn = -55.2 kJ mole-1


CH3COOH  +  NH4OH                       CH3COONH4  +  H2O    ΔHθn = -50.3 kJ mole-1

Conservation of energy and Hess's Law to calculate enthalpy changes.

Hess’s Law

Many reaction enthalpies cannot be determined experimentally in the laboratory e.g. the enthalpy of formation of ethanol.  It can be determined indirectly by applying Hess's law which states that the total enthalpy change accompanying a chemical change is independent of the route by which the chemical change takes place.

 

Consider the reaction:


           A  +  B            C  +  D

There may be more than one route for this reaction to take place.

                                             Route 1

 

                                                                    ΔH1                                                 

A  +  B                                                                                          C  +  D

                                       ΔH3

   ΔH

2             E  +  F                           G +  H                          ΔH4

                                             Route 2

By Hess’s Law

           Enthalpy change for Route 1  =  Enthalpy change for Route 2

                       ΔH1  = ΔH2  + ΔH3  + ΔH4

Hess’s Law can be used to find the enthalpy change for reactions that would be impossible to carry out in the laboratory.

e.g. the standard enthalpy of formation of ethyne is impossible to determine from practical measurements.


           2C (s)  +  H2 (g)                                         C2H2 (g)

Attempts to carry this out would result in the formation of a mixture of hydrocarbons. However the enthalpy can be calculated indirectly from enthalpies of combustion which can be measured accurately.

                                             ΔH1


           2C (s)  +  H2 (g)                                   C2H2 (g)

    ΔH2                    ΔH3           ΔH4    

 

2CO2 (g)   +  H2O (l)

By Hess’s Law:

ΔH1  + ΔH4  = ΔH2  + ΔH3                       [OR ΔH1  + ΔH4  - ΔH2  - ΔH3  =  0]

Where ΔH1 = enthalpy of formation of ethyne

ΔH2 = 2 x enthalpy of combustion of carbon 

ΔH3 = enthalpy of combustion of hydrogen

ΔH4 = enthalpy of combustion of ethyne

 

Worked Examples

1.<