R1.1.4 Standard enthalpy change

Understanding Standard Enthalpy Change and Heat Calculations

1. Picture yourself in a chemistry lab.
2. You're holding a beaker of water with a thermometer immersed in it.
3. You add a small piece of magnesium ribbon to hydrochloric acid, and suddenly, bubbles form, and the solution begins to warm up.
4. The thermometer shows a noticeable increase in temperature. What’s happening here?

This is a classic example of energy transfer during a chemical reaction, and the heat released can be analyzed using the concept of enthalpy change.

What is Standard Enthalpy Change ($\Delta H^{\ominus }$)?

Chemical reactions are accompanied by energy transfers between the system (reactants and products) and the surroundings.

When this measurement is conducted under standard conditions—defined as 298 K (25°C), 100 kPa, and reactants/products in their standard states—it is referred to as the standard enthalpy change($\Delta H^{\ominus }$).

• Exothermic reactions: Release heat into the surroundings, resulting in a negative $\Delta H$ (e.g., combustion of fuels).
• Endothermic reactions: Absorb heat from the surroundings, resulting in a positive $\Delta H$ (e.g., melting of ice).

Units of Enthalpy Change

• The standard unit for $\mathrm{\Delta }H$ is kilojoules per mole (kJ mol⁻¹).
• This unit reflects the energy change associated with one mole of reactant or product under standard conditions.

Calculating Heat Transfer $Q$

The heat transferred during a reaction can be calculated using the formula:

$$Q=mc\mathrm{\Delta }T$$

where:

• $Q$: Heat transferred (in joules, J)
• $m$: Mass of the substance being heated (in grams, g)
• $c$: Specific heat capacity of the substance (in J g⁻¹ K⁻¹)
• $\mathrm{\Delta }T$: Temperature change (in kelvin, K, or degrees Celsius, °C)

Breaking Down the Formula

1. Mass $m$: This is the mass of the substance (often water in calorimetry experiments) that absorbs or loses heat.
2. Specific Heat Capacity $c$: A property of the substance that indicates how much heat is required to raise the temperature of 1 g of the substance by 1 K. For water, $c=4.18\mathrm{J}{\mathrm{g}}^{-1}{\mathrm{K}}^{-1}$.
3. Temperature Change $\mathrm{\Delta }T$: Calculated as $\mathrm{\Delta }T=T_{\text{final}}-T_{\text{initial}}$.

Connecting Heat ($Q$) to Enthalpy Change ($\Delta H$)

The relationship between heat transfer $Q$ and enthalpy change $\mathrm{\Delta }H$ is given by:

$$\mathrm{\Delta }H=-\frac{Q}{n}$$

where:

• $\mathrm{\Delta }H$: Enthalpy change (in kJ mol⁻¹)
• $Q$: Heat released or absorbed (in joules or kilojoules)
• $n$: Number of moles of the limiting reactant (in moles)

Why the Negative Sign?

• The negative sign reflects the system’s perspective.
• If the system releases heat (exothermic reaction), $Q$ is positive, but $\mathrm{\Delta }H$ is negative.
• Conversely, if the system absorbs heat (endothermic reaction), $Q$ is negative, but $\mathrm{\Delta }H$ is positive.

Practical Applications of Enthalpy Calculations

1. Calorimetry Experiments

• In lab settings, calorimeters (e.g., coffee-cup calorimeters) measure heat changes during reactions.
• These experiments provide valuable insights into reaction energetics, aiding in the design of efficient chemical processes.

2. Industrial Processes

• Understanding enthalpy changes helps engineers optimize processes such as fuel combustion in power plants or ammonia production in the Haber process.

3. Environmental Impact

• Knowledge of enthalpy changes is crucial for assessing the energy efficiency and environmental impact of chemical reactions, such as the combustion of fossil fuels.

Reflection and Practice