Welcome to the World of Enthalpy!

In this chapter, we are going to explore why some chemical reactions get hot enough to cook your dinner while others get cold enough to freeze water. We call the study of these energy changes Thermodynamics. Specifically, we are looking at Enthalpy, which is just a fancy scientific word for the "heat content" of a system. By the end of these notes, you’ll be able to predict energy changes, draw diagrams to show them, and calculate exactly how much heat is moving around!

Prerequisite Concept: Remember that energy cannot be created or destroyed; it only moves from one place to another. In Chemistry, energy moves between the system (the chemicals) and the surroundings (everything else, like the water in a beaker or the air in the room).


1. Exothermic and Endothermic Reactions

Chemical reactions always involve a change in energy because bonds are being broken and new ones are being made. We use the symbol \(\Delta H\) (delta H) to represent the Enthalpy Change.

Exothermic Reactions (\(\Delta H\) is negative)

In an exothermic reaction, the chemicals lose heat energy to the surroundings. Analogy: Think of it like a bank account. If you spend money (energy), your balance goes down (negative change), but the person you paid (the surroundings) gets richer (warmer).

  • The temperature of the surroundings increases.
  • \(\Delta H\) is negative (e.g., \(-57 \text{ kJ mol}^{-1}\)).
  • Examples: Combustion (burning fuel), neutralisation, and respiration.

Endothermic Reactions (\(\Delta H\) is positive)

In an endothermic reaction, the chemicals absorb heat energy from the surroundings. Analogy: This is like depositing money into your bank account. Your balance goes up (positive change), but the money has to come from somewhere else!

  • The temperature of the surroundings decreases.
  • \(\Delta H\) is positive (e.g., \(+120 \text{ kJ mol}^{-1}\)).
  • Examples: Thermal decomposition (breaking down limestone) and photosynthesis.

Memory Aid: Exo sounds like Exit (Heat leaves). Endo sounds like Enter (Heat enters).

Key Takeaway: Exothermic reactions release heat (negative \(\Delta H\)), while endothermic reactions absorb heat (positive \(\Delta H\)).


2. Enthalpy Profile Diagrams

These diagrams are a visual "map" of the energy changes in a reaction. They show the enthalpy of the reactants versus the products.

Important Features:

  • The y-axis: Represents Enthalpy (\(H\)).
  • The x-axis: Represents the progress of the reaction.
  • Activation Energy (\(E_a\)): This is the "energy hill" the chemicals must climb to start the reaction. It is the minimum energy required for a reaction to take place.

Exothermic Profile

The products are lower than the reactants because energy was lost. The arrow for \(\Delta H\) points downwards.

Endothermic Profile

The products are higher than the reactants because energy was gained. The arrow for \(\Delta H\) points upwards.

Don't worry if this seems tricky at first! Just remember: if the line ends lower than it started, it's exothermic. If it ends higher, it's endothermic.


3. Standard Enthalpy Changes

To compare reactions fairly, chemists use Standard Conditions. This ensures everyone is measuring the same thing. Look for the "theta" symbol (\(^\ominus\)) to know if conditions are standard (e.g., \(\Delta H^\ominus\)).

Standard Conditions are:

  • Pressure: \(100 \text{ kPa}\) (roughly normal atmospheric pressure).
  • Temperature: \(298 \text{ K}\) (which is \(25^\circ \text{C}\)).
  • Concentration: \(1.0 \text{ mol dm}^{-3}\) (for solutions).
  • Standard State: The physical state (solid, liquid, or gas) a substance is in under these conditions.

Key Definitions you need to know:

1. Enthalpy change of reaction (\(\Delta_r H^\ominus\)): The enthalpy change that accompanies a reaction in the molar quantities shown in a chemical equation.

2. Enthalpy change of formation (\(\Delta_f H^\ominus\)): The enthalpy change when one mole of a compound is formed from its elements in their standard states. Note: The \(\Delta_f H^\ominus\) of any element is always \(0 \text{ kJ mol}^{-1}\).

3. Enthalpy change of combustion (\(\Delta_c H^\ominus\)): The enthalpy change when one mole of a substance reacts completely with oxygen.

4. Enthalpy change of neutralisation (\(\Delta_{neut} H^\ominus\)): The enthalpy change when an acid and a base react to form one mole of water.

Quick Review: Formation = forming 1 mole of product. Combustion = burning 1 mole of reactant. Neutralisation = forming 1 mole of water.


4. Measuring Enthalpy Changes Directly

We can calculate the energy change in a lab using a technique called calorimetry. We use a simple formula to find the heat energy (\(q\)) exchanged with the surroundings (usually water):

\(q = m \times c \times \Delta T\)

  • \(q\): Heat energy (in Joules, J).
  • \(m\): Mass of the surroundings being heated (usually the water or solution, in grams).
  • \(c\): Specific heat capacity (for water, this is always \(4.18 \text{ J g}^{-1} \text{ K}^{-1}\)).
  • \(\Delta T\): Change in temperature (final temp \(-\) initial temp).

Converting \(q\) to Enthalpy Change (\(\Delta H\)):

To find \(\Delta H\) in \(\text{kJ mol}^{-1}\), follow these steps:

  1. Calculate \(q\) using \(mc\Delta T\).
  2. Convert \(q\) from Joules to kiloJoules (divide by \(1000\)).
  3. Calculate the number of moles (\(n\)) of the chemical that reacted.
  4. Divide \(q\) by \(n\) (\(\Delta H = q / n\)).
  5. Important: Add a minus sign if the temperature went up (exothermic)!

Common Mistake to Avoid: When using \(q = mc\Delta T\), the mass (\(m\)) is the mass of the liquid you are heating, NOT the mass of the solid you might be adding to it.


5. Bond Enthalpies

Bonds act like rubber bands. Breaking them requires energy, and forming them releases energy.

  • Breaking bonds: Endothermic (requires energy).
  • Making bonds: Exothermic (releases energy).

Memory Aid: MEXO BENDO (Making = Exo / Breaking = Endo).

Average Bond Enthalpy

This is the energy needed to break one mole of a specified type of bond in a gaseous molecule. We use "average" because the exact energy depends on the environment of the bond.

Calculating \(\Delta H\) from Bond Enthalpies:

\(\Delta H = \sum(\text{bond enthalpies of reactants}) - \sum(\text{bond enthalpies of products})\)

Or simply: \(\Delta H = \text{Bonds Broken} - \text{Bonds Made}\)

Key Takeaway: If the energy released when making new bonds is greater than the energy used to break the old ones, the reaction is exothermic.


6. Hess' Law and Enthalpy Cycles

Sometimes we can't measure a reaction directly (maybe it's too dangerous or too slow). Hess' Law says that the total enthalpy change for a reaction is the same regardless of the route taken.

Type A: Using Enthalpy of Formation (\(\Delta_f H\))

In these cycles, the Elements are at the bottom. The arrows point UP from the elements to the reactants and products.

Equation: \(\Delta H = \sum \Delta_f H (\text{products}) - \sum \Delta_f H (\text{reactants})\)

Type B: Using Enthalpy of Combustion (\(\Delta_c H\))

In these cycles, the Combustion Products (\(\text{CO}_2\) and \(\text{H}_2\text{O}\)) are at the bottom. The arrows point DOWN from the reactants and products to the combustion products.

Equation: \(\Delta H = \sum \Delta_c H (\text{reactants}) - \sum \Delta_c H (\text{products})\)

Did you know? Hess' Law is just a specific version of the Law of Conservation of Energy. It's like saying if you travel from London to Manchester, the change in your elevation is the same whether you drive directly or take a detour through Birmingham!

Summary of Hess' Law:
If you have Formation data: Products - Reactants.
If you have Combustion data: Reactants - Products.


Quick Review Quiz

1. If a reaction feels cold, is it exothermic or endothermic? (Answer: Endothermic)

2. What is the value of \(\Delta_f H^\ominus\) for \(\text{O}_2(g)\)? (Answer: \(0 \text{ kJ mol}^{-1}\), because it is an element)

3. Does breaking a chemical bond release or require energy? (Answer: Requires energy - BENDO)