Energetic/Thermochemistry (IB)

Measuring Energy Changes

Energetics is the study of heat changes in chemical reactions. Heat is a form of energy, and in chemistry, we measure how much heat is absorbed or released during reactions. For example, when you burn wood, heat is released into the surroundings.

Temperature and Energy

• Temperature: Measures how hot or cold something is. It’s related to the average kinetic energy (energy of motion) of particles in a substance. For example, boiling water has higher kinetic energy than ice.

• Heat: Measures the total energy content of a substance. A beaker of boiling water has more heat energy than a few drops of boiling water, even though both are at the same temperature.

• First Law of Thermodynamics: Energy cannot be created or destroyed, only transformed. For instance, when you burn fuel, chemical energy is transformed into heat and light.

• Second Law of Thermodynamics: For a reaction to happen on its own (spontaneously), there must be an overall increase in entropy (disorder). For example, ice melting into water increases entropy because water molecules are more disordered.

• Energy: The ability to do work. Energy cannot be created or destroyed, only transformed (First Law of Thermodynamics).

Enthalpy (ΔH)

Enthalpy is the heat energy stored in a substance, mainly in its chemical bonds. We can’t measure the total enthalpy of a substance, but we can measure the change in enthalpy (ΔH) during a reaction.

• Units: kJ/mol.

• System vs. Surroundings:

  • System: The reaction itself (e.g., a test tube with reactants).

  • Surroundings: Everything outside the system (e.g., the air around the test tube).

Standard Conditions

• Pressure: 100 kPa.

• Concentration: 1 mol/dm³ for solutions.

• Temperature: Usually 298 K (25°C).

• Standard State: Each substance is in its most stable form (solid, liquid, or gas).

Types of Standard Enthalpy Changes

• Standard Enthalpy Change of Reaction (ΔHᵣ): Enthalpy change when reactants react to form products under standard conditions.

• Standard Enthalpy Change of Formation (ΔHᶠ): Enthalpy change when 1 mole of a compound is formed from its elements under standard conditions.

• Standard Enthalpy Change of Combustion (ΔH_c): Enthalpy change when 1 mole of a substance is burned in excess oxygen under standard conditions.

• Standard Enthalpy Change of Neutralization (ΔH_neut): Enthalpy change when 1 mole of water is formed by reacting an acid and alkali under standard conditions.

Exothermic vs. Endothermic Reactions

• Exothermic Reactions: Release heat to the surroundings. ΔH is negative. Example: Combustion of propane:

  • C3H8+5O2→3CO2+4H2O(ΔH=−2220 kJ/mol)

  • The products are more stable than the reactants.

• Endothermic Reactions: Absorb heat from the surroundings. ΔH is positive. Example: Photosynthesis:

  • 6CO2+6H2O→C6H12O6+6O2(ΔH=+2800 kJ/mol)

  • The products are less stable than the reactants.

Measuring Enthalpy Changes

To measure enthalpy changes experimentally, we use the formula:

Q=mcΔT

• Q: Heat energy (in kJ).

• m: Mass of water (in kg).

• c: Specific heat capacity of water (4.18 kJ/kg·K).

• ΔT: Change in temperature (in K).

Then, the enthalpy change (ΔH) is calculated as:

ΔH= Q/n

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• n: Number of moles of the limiting reactant.

Calorimetry

Measuring Enthalpy Changes

• Calorimeter: A device used to measure heat changes in reactions. Can be made from a polystyrene cup, vacuum flask, or metal can.

• Specific Heat Capacity (c): The energy needed to raise the temperature of 1 g of a substance by 1 K. For water, c = 4.18 J g⁻¹ K⁻¹.

• Equation: q=m×c×ΔT

Worked Example

• Problem: 0.01 mol of propan-1-ol was burned, heating 250 g of water from 298 K to 310 K. Calculate the enthalpy of combustion.

• Solution:

  • q=250×4.18×12=12,540 J

  • ΔH=q/n=12,54/00.01=1,254,000 J/mol=−1,254 kJ/mol

Calorimetry Experiments

Enthalpy Changes in Solution

• Principle: Reactants are mixed in solution, and the temperature change is measured.

• Assumptions:

  • Specific heat capacity of the solution = 4.18 J g⁻¹ K⁻¹.

  • Density of the solution = 1 g/cm³.

  • Heat losses are negligible.

Temperature Correction Graphs

• Purpose: To account for heat loss during slow reactions.

• Method:

  1. Record temperature before adding reactants.

  2. Add reactants and continue recording temperature.

  3. Extrapolate the cooling section of the graph to find the maximum temperature change.

Enthalpy of Combustion Experiments

• Principle: Heat from combustion is used to heat water.

• Sources of Error:

  • Heat loss to surroundings.

  • Incomplete combustion.

Worked Example

• Problem: 1.023 g of propan-1-ol was burned, heating 200 g of water by 30°C. Calculate the enthalpy of combustion.

• Solution:

  • q=200×4.18×30 = 25,080 J

• ΔH=q/n=25,080/0.01702=1,473,560 J/mol=−1,474 kJ/mol

Hess’s Law

Hess’s Law states that the total enthalpy change for a reaction is the same, no matter the route taken. For example, if you want to find the enthalpy change for:

A→C

You can use intermediate steps:

A→B(ΔH1)B→C(ΔH2)

Then:

ΔHtotal=ΔH1+ΔH2

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Bond Enthalpies

• Bond Breaking: Requires energy (endothermic). Example: Breaking H-H bonds in hydrogen gas.

• Bond Making: Releases energy (exothermic). Example: Forming H-O bonds in water.

The enthalpy change for a reaction can be calculated using bond enthalpies:

ΔH=Bond Enthalpies of Reactants−Bond Enthalpies of Products

Energy Cycles

Born-Haber Cycle

The Born-Haber Cycle is used to calculate the lattice enthalpy of ionic compounds. It breaks down the formation of an ionic compound into steps:

1. Atomization: Convert solid metal to gaseous atoms.

2. Ionization: Remove electrons from metal atoms to form cations.

3. Electron Affinity: Add electrons to non-metal atoms to form anions.

4. Lattice Formation: Combine gaseous ions to form the solid lattice.

For example, to calculate the lattice enthalpy of LiF:

ΔHlat=ΔHatom+ΔHIE+ΔHEA−ΔHf

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Entropy and Spontaneity

Entropy (S)

Entropy measures the disorder of a system. The more ways energy can be distributed, the higher the entropy. For example:

• Solids have low entropy (ordered).

• Gases have high entropy (disordered).

The change in entropy (ΔS) is calculated as:

ΔS=Sproducts−Sreactants

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Gibbs Free Energy (ΔG)

Gibbs free energy combines enthalpy and entropy to predict whether a reaction is spontaneous:

ΔG=ΔH−TΔS

ΔG=ΔH−TΔS

• If ΔG is negative, the reaction is spontaneous.

• If ΔG is positive, the reaction is not spontaneous.

Example: At high temperatures, even endothermic reactions (ΔH > 0) can be spontaneous if ΔS is large enough.