Enthalpy Changes

Section 1 of 14Enthalpy and enthalpy change

• Enthalpy = the total chemical energy (heat content) inside a substance
  • The enthalpy change is ΔH, where Δ means change and H is enthalpy.
  • ΔH can be positive or negative.

Section 2 of 14Exothermic and endothermic reactions

Feature	Exothermic	Endothermic
Energy of products	Less than reactants	More than reactants
Heat transfer	Given out to the surroundings	Absorbed from the surroundings
Surroundings temperature	Increases	Decreases
Energy of the system	Decreases	Increases
ΔH	Negative	Positive

• Feasibility
  • an exothermic reaction is thermodynamically possible, but if the rate is too slow it may not occur
  • Such a reaction is kinetically controlled.

Section 3 of 14Enthalpy profile diagrams

Features of an energy level diagram

• The energy level of the reactants.
• The transition state, an unstable intermediate that cannot be isolated and is higher in energy than reactants and products (the term itself need not be labelled).
• The energy level of the products.
• The activation energy, E_a.
• The enthalpy change ΔH.

Activation energy

Activation energy the minimum energy needed for reactant molecules to collide successfully and start the reaction

Enthalpy change

Enthalpy change the overall energy taken in from or given out to the surroundings, i.e. the energy difference from reactants to products

1

Draw labelled axes for energy and progress of reaction.

2

Place reactants higher than products for an exothermic reaction, and lower for an endothermic reaction.

3

Draw the curve with a hump between them.

4

Draw arrows for E_a (reactants to the top of the hump) and ΔH (reactants to products).

• Reverse reaction
  • E_a(forward) = ΔH + E_a(reverse) for an endothermic forward reaction
  • ΔH+70 and E_a(reverse) = +20 kJ mol⁻¹ give E_a(forward) = +90 kJ mol⁻¹.

Section 4 of 14Standard conditions

Standard conditions

• Standard conditions used so enthalpy changes can be compared
• A pressure of 100 kPa (older questions may use 101 kPa; the exact figure is 101 325 Pa).
• A temperature of 298 K (25 °C).
• Each substance in its standard physical state.

Symbol

• Symbol a standard enthalpy change is written ΔH with the standard symbol (ΔH^⊗)

Section 5 of 14Standard enthalpy definitions

Enthalpy change	Definition	Sign
Reaction, ΔHr	Enthalpy change when the reactants in the stoichiometric equation react to form the products, all in standard states	Either
Formation, ΔHf	Enthalpy change when one mole of a compound is formed from its elements in their standard states	Either
Combustion, ΔHc	Enthalpy change when one mole of a substance in its standard state is burnt in excess oxygen	Exothermic
Neutralisation, ΔHneut	Enthalpy change when one mole of water is formed by reacting an acid and an alkali	Exothermic

Standard enthalpy definitions

Elements

• ΔH_f of an element in its standard state is zero, e.g. O₂(g) is 0 kJ mol⁻¹

Scaling

• Scaling multiply ΔH by the number of moles in the equation
• H₂ + ½O₂ → H₂O is −286 kJ mol⁻¹, so 2H₂ + O₂ → 2H₂O is −572 kJ mol⁻¹.

Identifying the type

• C(graphite) + O₂ → CO₂ is both ΔH_f and ΔH_c.
• HCl + NaOH → NaCl + H₂O is ΔH_neut because one mole of water forms.

Section 6 of 14Calorimetry

Calorimetry

Calorimetry

Calorimetry measuring enthalpy changes in chemical reactions

• A simple calorimeter can be a polystyrene cup, vacuum flask or metal can.

Specific heat capacity, c

Specific heat capacity, c the energy needed to raise the temperature of 1 g of a substance by 1 °C

• For water it is 4.18 J g⁻¹ K⁻¹.

Energy transferred

• Energy transferred q = m × c × ΔT

  What the symbols mean: E, Q or q = energy transferred (J) · m = mass · c = specific heat capacity (J kg⁻¹ °C⁻¹; water 4.18 J g⁻¹ °C⁻¹) · ΔT or Δθ = temperature change (°C or K)
  • (kg; g in chemistry calorimetry)
• Aqueous solutions are assumed to be mostly water, so use the m and c of water.

Enthalpy change per mole

• ΔH = q ÷ n = ÷ n

  What the symbols mean: ΔH = enthalpy change (kJ mol⁻¹) · q = heat energy measured (J, divide by 1000 for kJ) · n = amount of the limiting reactant (mol); negative sign = exothermic
  • (m × c × ΔT)
• A temperature rise gives a negative ΔH (exothermic, e.g. combustion).
• A temperature fall gives a positive ΔH (endothermic).

Example

• 2.50 g methane heats 500 g water from 25 to 68 °C, 30% of the energy released
• q = 500 × 4.18 × 43 = 89 870 J, which is 30% of the total, so total = 299 567 J.
• Per gram of methane this is 120 kJ g⁻¹.

Section 7 of 14Bond enthalpy

Bond enthalpy, E

Bond enthalpy, E the energy required to break one mole of a specific covalent bond in gaseous molecules

• It is also called bond energy or bond dissociation energy.
• E(H–H) is the bond enthalpy of single bonds between two hydrogen atoms.

Average bond enthalpy

• Average bond enthalpy bond enthalpies depend on the other atoms in the molecule, so an average over different environments is used
• It is found using enthalpy cycles because bond enthalpies cannot be measured directly.
• Data tables list average values, which may differ from the actual value for a molecule.

Section 8 of 14Calculating enthalpy change from bond enthalpies

1Use

• Use to find ΔH_r when it cannot be measured experimentally

2Energy flow

• Breaking bonds is endothermic, so its ΔH is positive.
• Making bonds is exothermic, so its ΔH is negative.
• The balance decides whether the overall reaction is exothermic or endothermic.

1

Work out the bonds broken and formed from the equation.

2

Set out a balance sheet of each bond type multiplied by its number.

3

Add the bonds broken as positive values and the bonds formed as negative values.

4

ΔH_r = Σ(bonds broken) + Σ(bonds formed).

• Example
  • Haber process N₂ + 3H₂ ⇌ 2NH₃
  • Bonds broken total +2253 kJ mol⁻¹ and bonds formed total −2346 kJ mol⁻¹.
  • ΔH_r = −93 kJ mol⁻¹.

Section 9 of 14Bond enthalpy calculation for combustion

Definition requirement

• Definition requirement the enthalpy of combustion is for one mole of substance reacting in excess oxygen to form water and carbon dioxide
• The equation must be scaled so only one mole of fuel reacts.

Example

• Example ethyne, H–C≡C–H + 2½O=O → H–O–H + 2O=C=O
• Bonds broken total +2912 kJ mol⁻¹ and bonds formed total −4142 kJ mol⁻¹.
• ΔH_c = −1230 kJ mol⁻¹.

Section 10 of 14Hess law

Hess law

Hess law the total enthalpy change of a reaction is independent of the route taken, provided the initial and final conditions are the same

• The enthalpy change of the direct route equals that of any indirect route.
• The principle matters more than the exact wording.

Use

• Use to find enthalpy changes that cannot be measured by calorimetry
• Δ_fH of propane, 3C + 4H₂ → C₃H₈, cannot be found directly because carbon and hydrogen do not react under standard conditions.

Section 11 of 14Cycles using enthalpies of formation

1Relationship

• Relationship products form directly from elements (ΔH₂) or indirectly via the reactants (ΔH₁ + ΔH_r)
• ΔH₂ = ΔH₁ + ΔH_r, so ΔH_r = ΔH₂ − ΔH₁.

2Rule

• Δ_rH = ΣΔH_f(products) − ΣΔH_f(reactants)

1

Write the equation for the reaction.

2

Write the elements, with correct moles and state symbols, underneath.

3

Draw upward arrows from the elements to each compound.

4

Write the ΔH_f values on the arrows, multiplied by the moles.

5

Go from reactants to products via the elements, reversing the sign when travelling against an arrow.

• Example
  • NH₄NO₃ + ½C → N₂ + 2H₂O + ½CO₂
  • ΔH_r = +365 − 484 − 197 = −316 kJ mol⁻¹.
  • No arrows are needed for C and N₂ because ΔH_f of an element is 0.

Section 12 of 14Cycles using enthalpies of combustion

1Rule

• Δ_rH = ΣΔH_c(reactants) − ΣΔH_c(products)

2Use

• Use enthalpies of formation are hard to measure but many enthalpies of combustion can be measured, so combustion data give ΔH_f

1

Write the equation for the formation of the compound.

2

Write the combustion products underneath.

3

Draw downward arrows from each substance to its combustion products.

4

Write the ΔH_c values on the arrows, multiplied by the moles.

5

Go from reactants to products, reversing the sign when travelling against an arrow.

1Example

• Example propanone, 3C + 3H₂ + ½O₂ → CH₃COCH₃
• ΔH_f = −1182 − 858 + 1821 = −219 kJ mol⁻¹.
• The sign of −1821 is reversed because the cycle runs against its arrow.

2Care

• Care keep each value in brackets so a minus sign is not lost, and balance the atoms of every element in the cycle

Section 13 of 14Key terms

Term	Meaning
Enthalpy change	The heat energy change of a reaction at constant pressure, ΔH
Activation energy	The minimum energy needed for a successful collision to start a reaction
Standard enthalpy of formation	Enthalpy change when one mole of a compound forms from its elements in standard states
Standard enthalpy of combustion	Enthalpy change when one mole of a substance is burnt completely in excess oxygen
Specific heat capacity	Energy needed to raise the temperature of 1 g of a substance by 1 K
Bond enthalpy	Energy needed to break one mole of a specific covalent bond in gaseous molecules

Term	Meaning
Hess law	Total enthalpy change is independent of the route taken
Enthalpy cycle	A diagram linking direct and indirect routes between reactants and products
Standard state	The physical state of a substance at 100 kPa and 298 K
Balance sheet	A table of bonds broken (positive) and bonds formed (negative) used to calculate ΔH

Section 14 of 14Exam technique

1

Writing a positive ΔH for an exothermic reaction — exothermic reactions have negative ΔH

2

Using the mass of the solid instead of the mass of water in q = mcΔT — use the water (solution) mass

3

Forgetting to divide by moles to convert q to ΔH per mole

4

Giving bonds formed a positive sign in a bond enthalpy balance — bonds formed are negative

1

Using Σ(reactants) − Σ(products) for formation data — it is products minus reactants for ΔH_f

2

Reversing the combustion rule — it is reactants minus products for ΔH_c

3

Not changing the sign when a cycle runs against an arrow

4

Forgetting to multiply ΔH by the number of moles in the balanced equation