What is chemical energy?

• Chemical energy is a form of potential energy which is stored in chemical bonds.
• Chemical bonds are the attractive forces that bind atoms together.
• As a reaction takes place, bonds break in the reactants and new bonds are formed in the products.
• The difference in energies of the bonds between the atoms of reactants and products is called chemical energy.

Enthalpy, H

• Enthalpy, H, is the heat content that is stored in a chemical system, as reactants or products.
• It is impossible to measure the enthalpy content of a system directly but we can measure the differences in enthalpy contents.
• The difference in enthalpy contents is the energy either given out or absorbed.

• The energy exchanged with the surroundings can be given out as light, sound, electrical but is more usually heat.
• This change in enthalpy content is called enthalpy change.

Conservation of energy (First Law of thermodynamics)

Energy cannot be created or destroyed just transferred between the system and the surroundings

• The surroundings, water, tube, air either gains the energy given out by the system (hot)
• Or the surroundings, water, tube, air loses energy to the system (cold) - like ethanol evaporating from your hand
• This means that:

• What your thermometer is measuring is the surroundings.  That is - the water, tube, air the energy is transferred from / to.

Enthalpy change, DH:

• This is the difference between the enthalpy contents of the products and the reactants (at constant pressure):

                DH    =    H_products    -    H_reactants

• It is rare for enthalpy contents of products and reactants to be exactly equal.
• This means that there is almost always a transfer of energy between the system and the surroundings.

Exothermic reactions:

• This means that the enthalpy content of the products is smaller than the reactants.
• The excess energy is transferred from the system to the surroundings - hot

                                H_products    <    H_reactants

	DH	=	Hproducts	-	Hreactants
	DH	=	Small	-	Large
	DH	=	Negative value
Energy is released from the system to the surroundings

Endothermic reactions:

• This means that the enthalpy content of the products is greater than the reactants.
• The excess energy is transferred to system from the surroundings - cold

                                H_products    >    H_reactants

	DH	=	Hproducts	-	Hreactants
	DH	=	Large	-	Small
	DH	=	Positive value
Energy is transferred to the system from surroundings

Questions 1 - 2  P 185

Exothermic and endothermic reactions

1)  Oxidation of fuels:

• The most common example is the oxidation of methane forming carbon dioxide and water:

• The negative sign means that the reaction is exothermic.

• This means that H_products    <    H_reactants

• The units tell you that 890 Kj of energy is given out per mole of methane.

• Changes in enthalpy are given in molar quantities:

Environmental impact - Example:

• A small car travelling 2 miles uses 200 - 250g of petrol, C₈H₁₈, DH = -5470 KJ mol^-1. 

• How much energy is needed and how much CO₂ is produced?

2)  Respiration:

• An important exothermic reaction providing energy for all living things:

Endothermic reactions:

1)  Photosynthesis:

• An important endothermic reaction without which there would be no life.

• This reaction makes 'food', the starting point for all foods, and oxygen for respiration.

• Energy from the sun provides the energy to transfer from surroundings to system.

• The reverse reaction of respiration.

2)  Thermal decomposition of limestone:

• Limestone contains calcium carbonate:

• Calcium oxide is commonly known as lime used in cement and used to treat soils by farmers.

• If water is added to lime, calcium oxide, heat is given off - exothermic reaction.

Questions 1 - 2  P187

Enthalpy profile diagrams

Simple enthalpy profile diagrams:

• Reactions and their enthalpy changes have been shown as simple enthalpy profile diagrams:

Exothermic reactions: Where Hproducts  <  Hreactants DH = negative	Endothermic reactions: Where Hreactants  >  Hproducts DH = positive

Activation energy: 

• The simple energy profile diagrams assumes that as soon as the reactants come in contact with each other a reaction takes place.

• Most reactions do not occur 'spontaneously' but need a little bit of energy to get them going, a spark is needed to set gas alight.

• This 'bit of energy' is called the activation energy.  It is the energy required to break the bonds in the reactants.

Exothermic reactions:	Endothermic reactions:
Even though the products are lower in energy than the reactants, a small amount of energy is needed to break the reactant bonds, the activation energy. We say the reaction has to overcome the energy barrier. After that, the exothermic nature of the reaction is enough to break more reactant bonds. A way of thinking of this is like a cyclist riding a bike up a small hill to gain a large amount of 'free wheeling' or energy.	The products are higher in energy than the reactants, a small amount of energy is still needed to break the reactant bonds, the activation energy. The reaction has to overcome the energy barrier. This time there is no excess energy to break more reactant bonds, a sustained amount of energy needs to be continually supplied to keep the reaction going. This time the cyclist rides a bike up a large hill to gain a small amount of 'free wheeling' or energy.

Questions  1-2  P189

Standard enthalpy changes

Standards:

• Enthalpy changes for reactions will vary slightly depending upon the conditions under which the reaction is carried out.

• All data books have to have the same value for enthalpy changes.

• Chemists use standard conditions to ensure that all reactions and corresponding enthalpy changes are carried out under the same conditions.

• They are as close to normal lab conditions as possible.

Standard conditions:

                      Pressure at 100 kPa (which is 1 atmosphere)

                      Temperature at 298K (25^oC)

                      1 Mole or 1 Molar solutions

                      Normal physical states at standard temperature and pressure (above conditions)

                      A standard enthalpy change is shown by  -   DH ^q

                       H - Enthalpy

                       D - Change in

               q symbol represents standard conditions

Standard states:

• Standard enthalpy changes must have substances in their standard states under these standard conditions:

Standard enthalpy changes:

• There are 3 standard enthalpy changes that you need to know:

1)  Standard enthalpy change of reaction, DH_r^q

• For this enthalpy change we need a reaction to refer to.

• A value for the enthalpy change which is true for the molar quantities stated in the reaction:

For the reaction:

         Fe₂O_3(s)   +    2Al_(s)     à      2Fe_(s)    +       Al₂O_3(s)             DH^q_r   =  -851.50 KJ Mol^-1

• For the same reaction with different stoichiometry:

   1/2Fe₂O_3(s)    +    Al_(s)        à      Fe_(s)       +       1/2Al₂O_3(s)      DH^q_r   =  -425.75 KJ Mol^-1

• Note that the reaction must be written in order for the information to be correct.  DH^q_r  is half the value if the reaction is balanced using halves.

2)  Standard enthalpy of combustion – DH^q_c

• Combustion reactions are so common that they have a standard enthalpy change of their own:

Definition:  The enthalpy change that occurs when 1 mole of a substance reacts completely with oxygen under standard conditions.  All reactants and products are in their standard states.

DH^q_c for ethane:

• DH^q_c refers to the complete combustion of 1 mole of ethane.

• This conforms with the definition -' when 1 mole of a substance reacts completely with oxygen'

3)  Standard enthalpy change of formation, DH^q_f

• Formation reactions are also so common that they have a standard enthalpy change of their own:

Is the enthalpy change that takes place when 1 mole of a compound is formed from its constituent elements in their standard states under the standard conditions.

DH^q_f  for water:

   H_2(g)            +          1/2O_2(g)           à         H₂O_(l)              DH^q_f  =  -286 KJ Mol^-1

• This refers to the Standard enthalpy change of formation of 1 mole of water.

• This conforms with the definition -' when 1 mole of a compound is formed from its elements'

A problem:

• If we are forming an element H₂, there is no chemical change.

• We say that elements in their standard states under standard conditions have a standard enthalpy of 0 KJ Mol^-1

Questions  1 - 3   P191

Determination of enthalpy changes

• Enthalpy content of reactants and products cannot be measured directly.

• You can measure the enthalpy change between the reactants and products - enthalpy change for a reaction, DH_r

Remember:

• The enthalpy change for any reaction, DH_r is measured in KJ Mol^-1, which is energy per mole

• This means we need to know 2 things

i)                    The energy change - for which we use temperature change

ii)                   The amounts in moles of the limiting reagent that reacts

i)  The energy change:

• To determine the energy change we use the formula:

                                                  Q = mcDT

                                                          1000

                                                  Q       – quantity of energy exchanged J, the 1000 converts to kJ

                                                  m      – mass of the water g (ie cm³ as  the density if water is 1 gcm^-3)

                                                  c       – specific heat capacity j g^-1 K^-1

                                                  DT    – rise in temperature K, (T_initial  -  T_final)

• Convert the energy calculated into KJ (divide by 1000).

ii)  Calculate the number of moles used:

                                                  No Moles = Mass         or        c  x  V

                                                                           Mr                       

iii)  Calculate the amount of energy exchanged per mole , this is the enthalpy:

                                                              Enthalpy = Energy

                                                                                    Moles

• Check you have the sign correct:

                                            (-)ve for exothermic reactions

                                            (+)ve for endothermic reactions

Direct determination of enthalpy changes:

Most reactions involve adding one reactant to another and measuring the temperature rise. This is called direct determination and is carried out in a calorimeter. A calorimeter is an insulated reaction vessel which minimises heat exchange to the air as the surroundings we measure is the water / solvent. The simplest calorimeter is a polystyrene cup but they can get quite complex.

Example:

Excess Mg is added to 100cm³ of 2.00 Mol dm^-3 CuSO₄, the temperature rose from 20.0^oC to 65^oC.

i)  The energy change:

ii)  Calculate the number of moles used:

                                                  No Moles = c  x  V

                                                  No Moles = 2  x  0.100

                                                  No Moles =  0.2

iii)  Calculate the amount of energy exchanged per mole , this is the enthalpy:

                                                              Enthalpy = Energy

                                                                                    Moles

                                                              Enthalpy = - 18.81

                                                                                       0.2

                                                              Enthalpy = - 94.05 Kj Mol^-1

Finally write the equation with the enthalpy change:

Questions  1 - 2   P193

Enthalpy change of combustion

• This is any reaction involving oxygen forming oxides:

Experimental determination of DH_c

• The enthalpy change of combustion, DH_c is also measured in KJ Mol^-1, which is energy per mole

• This means we need to know 2 things

i)                    The energy change - for which we use temperature change when heating water

ii)                   The amounts in moles of fuel used

To calculate DHc, measure a known mass (volume) of water - for Q = MCDT. Take the initial temperature of the water. Weigh the burner before heating. Heat to raise the temperature by about 10oC. Take a final temperature. Reweigh the burner to get a mass of fuel used.

Example:

1.5g of propan-1-ol heated 250cm³ of water by 45^oC.

i)  The energy change:

ii)  Calculate the number of moles used:

                                                  No Moles = Mass

                                                                           Mr

                                                  No Moles = 1.5

                                                                         60

                                                  No Moles =  0.025

iii)  Calculate the amount of energy exchanged per mole , this is the enthalpy:

                                                              Enthalpy = Energy

                                                                                    Moles

                                                              Enthalpy = - 47.025

                                                                                       0.025

                                                              Enthalpy = - 1881 Kj Mol^-1

Comparison of experimental value with standard enthalpy change:

Errors:

• Incomplete combustion of the fuel - less heat energy given out

• Heat loss to the surroundings - less heat energy measured

Improvements:

Use a Bomb calorimeter (left). This apparatus reduces heat loss s the water is insulated from the surroundings. It is burnt in oxygen to ensure complete combustion.

Questions  1-2  P195

Bond enthalpies

Bond enthalpy:

    Is the enthalpy change that takes place when breaking by homolytic fission1 mole of a given bond in the molecules of a gaseous species

• To summarise - it is the energy required to break 1 mole of bonds in the gaseous state.

• Some examples:

• These bonds enthalpies can only exist in these examples but some bonds can exist in very different molecules.

• In cases like this we use the Average bond enthaply

Average bond enthalpy:

• Bonds like C - H exist in all hydrocarbons but their bond enthalpies will vary depending on their environment:

• The C - H bond enthalpy will be different in an alkane than the C - H bond enthalpy in an aldehyde where the C atom is adjacent to an oxygen atom.

• For bond enthalpies that are common, we use average bond enthalpies:

Breaking and making bonds: 

• In this and any reaction, reactant bonds will be broken and new product bonds will be formed.

• When bonds are broken energy is required making it an endothermic process, the activation energy.

• When new bonds are formed energy is release making it an exothermic process.

• If the energy released when new bonds form making the products is greater than the energy needed to break the bonds of the reactants, the reaction is exothermic.

• If the energy released when new bonds form making the products is less than the energy needed to break the bonds of the reactants, the reaction is endothermic.

• Strong bonds require lots of energy to break but they also release lots of energy when they form:

Using bond enthalpies to determine enthalpy changes:

• The symbol   S   is used for 'sum of' or adding together

Formula:

Worked example:

	à		à
S  (Bond enthalpies of bonds broken)				S  (Bond enthalpies of bonds formed)

Questions 1-2  P197

Enthalpy changes from DH_c^q - Hess's law

Measuring enthalpy changes indirectly:

• Sometimes it is not possible to measure an enthalpy change directly.

• This may be due to:

                High activation energy

                Slow rate of reaction

                More than one reaction occurring at the same time

• Hess's law allows us to work out enthalpy changes that are not possible to measure:

Definition:  Hess’s Law 

The total enthalpy change accompanying a chemical change is independent of the route by which the chemical change takes place provided the initial and final conditions are the same.

• If this was not true, it could be possible to gain energy via route A instead of going via route B.  This would break the 1^st law of thermodynamics.

• It is found that reactants can be converted into the same products by more than 1 route.

• The total energy for each route is the same:  These are called Enthalpy cycles.

Using Hess’s law:  Enthalpy cycles

• Enthalpy cycles are used to determine enthalpy changes for reactions that are not easily measured directly.

• This can be done in 2 ways:

1.      Enthalpy changes of combustion

2.      Enthalpy changes of formation

1  Calculating DH^q  from enthalpy changes of combustion: 

• Consider the enthalpy change below:

                                  3C_(s)                +          4H_2(g)         à            C₃H_8(g)

• Combustion reactions can be used to find this enthalpy change as carbon hydrogen and propane all burn in oxygen.

• This means that their enthalpy changes of combustion can all be measured

• Constructing an enthalpy cycle:

• we use the opposite sign for - 2219 as our route, Route 2 goes the opposite way to the enthalpy change

            Route 2 = - 2326  +  + 2219

            Route 2 = - 107 Kj Mol^-1

Remember

            Route 1 = Route 2

            Route 1 = - 107 Kj Mol^-1

Using the formula:

Question 1  P199

Enthalpy changes from DH^q_f - Hess's law

2)  Using enthalpy change of formation   

• As with enthalpy changes of combustion, enthalpy changes of formation can be used in Hess's cycles to work out enthalpy changes.

• Consider the reaction:

                                  2SO_2(g)           +          O_2(g)         à            2SO_3(g)

• The enthalpy cycle would be:

• we use the opposite sign for - 594 = + 594 as our route, Route 2 goes the opposite way to the enthalpy change

            Route 2 = + 594  +  - 882

            Route 2 = - 288 Kj Mol^-1

Remember

            Route 1 = Route 2

            Route 1 = - 288 Kj Mol^-1

Using the formula:

Other enthalpy cycles:

• As long as there is a link between reactants and products, Hess's law can be applied and enthalpy cycles can be constructed.

• Follow the principles as outlined above and you won't go wrong.

Alternative approach using bond enthalpies:

Review:  Using the Formula:

Worked example:

	à		à
S  (Bond enthalpies of bonds broken)				S  (Bond enthalpies of bonds formed)

Alternative approach: 

Questions 1-2  P197

Summary of enthalpy cycles

Step 1:-    Write a balanced chemical equation for the reaction.

Step 2:-    Construct the Enthalpy cycle.

Step 3:-    Decide on your routes and draw them on the cycle

Step 4:-    Write in the DH^q  for each compound / element next to the arrows.

Step 5:-    Look up the values of each DH^q  and write them in.  Add them up for each route.

Step 6:-    Write out Hess’s law – Route 1  =  Route 2

Step 7:-    Put in your numbers.

Step 8:-    Calculate DH^q

Question 1-2    P201  /  1-7  P215  /  1,3-5  P217