Enthalpy, enthalpy change, and heat of reaction

Understanding chemistry involves not just looking at how substances react but also understanding the energy changes involved. When substances undergo chemical reactions, either heat is absorbed from the environment or released into the environment. This leads us to an important concept in thermochemistry called enthalpy. In this detailed explanation, we will understand enthalpy, enthalpy changes, and heat of reaction in a detailed manner that is easy to understand.

What is enthalpy?

Enthalpy (symbol: H) is the measure of the total energy of a thermodynamic system. It includes internal energy, which is the energy needed to create the system, and the energy needed to make room for it by displacing its environment and establishing its volume and pressure. In simple terms, enthalpy is related to the heat energy in chemical reactions taking place at constant pressure.

 H = U + PV 

In the above formula, H represents enthalpy, U is internal energy, P stands for pressure, and V is volume. Enthalpy itself is not that interesting; it provides valuable information about the change in enthalpy during the reaction.

Enthalpy change

The enthalpy change, expressed as ΔH, is the difference in enthalpy between the products and reactants of a chemical reaction. It is important for predicting whether a reaction will absorb or release heat. Enthalpy changes can be classified into two types:

• Exothermic reactions: These reactions release heat to the surrounding environment. For exothermic reactions the enthalpy change is negative ( ΔH < 0 ).
• Endothermic reactions: These reactions absorb heat from the surrounding. For endothermic reactions the enthalpy change is positive ( ΔH > 0 ).

Let us explore this concept with an example:

A + B → C + heat (exothermic reaction)
Heat + A + B → C (endothermic reaction)
    

In an exothermic reaction, heat is produced and given off, so ΔH is negative. In contrast, an endothermic reaction requires heat to proceed, which means ΔH is positive.

Understanding heat of reaction

The heat of reaction, also called the enthalpy change of reaction, is the enthalpy change when a reaction occurs. It is usually measured at standard conditions, represented as ΔH° where the degree sign indicates standard state conditions. The heat of reaction provides useful information about the energy changes in a reaction.

Standard enthalpy of formation

This is the enthalpy change when one mole of a compound is formed from its elements in their standard states. The formula is expressed as:

 ΔH°f 

For example, the formation of water from hydrogen and oxygen can be expressed as follows:

H₂(g) + ½O₂(g) → H₂O(l) ΔH°f = -285.8 kJ/mol
    

This shows that when one mole of water is formed, 285.8 kJ of energy is released.

Standard enthalpy of combustion

The standard enthalpy of combustion is the heat change that occurs when one mole of a substance reacts completely with oxygen under standard conditions. The combustion of methane is a classic example:

CH₄(g) + 2O₂(g) → CO₂(g) + 2H₂O(l) ΔH°c = -890.3 kJ/mol
    

This represents the energy released by the combustion of one mole of methane.

Visual examples with chemical equations and energy diagrams

The energy diagram above represents an exothermic reaction. It shows two reactants A + B being converted into products C + D with the release of energy. The difference in height between the levels of the reactants and products represents the energy released, which is represented by the downward slope.

This energy diagram shows an endothermic reaction. Here, the reactants absorb energy E + F and are converted into products G + H The upward slope indicates absorption of energy and the products are located at a higher energy level.

Factors affecting enthalpy change

Many factors affect the magnitude of the enthalpy change, such as:

1. Nature of reactants and products: Different substances have different internal energy, which affects the enthalpy change.
2. Physical state: The state of the reactants and products (solid, liquid, gas) affects the amount of energy change.
3. Bond energy: Enthalpy changes are also affected by the energy required to break bonds and the energy released when new bonds are formed.

For reactions involving bond breaking and formation, we can use bond enthalpy:

 ΔH = ∑(bond energy of reactants) – ∑(bond energy of products) 

Enthalpy calculation examples

Let us calculate the enthalpy change for some reactions using these principles. Consider:

Reaction: C(s) + O₂(g) → CO₂(g)
Given: ΔH°f [C(s)] = 0 kJ/mol, ΔH°f [O₂(g)] = 0 kJ/mol, ΔH°f [CO₂(g)] = -393.5 kJ/mol

ΔH = ΔH°f [products] - ΔH°f [reactants]
ΔH = [-393.5 kJ/mol] - [0 + 0]
ΔH = -393.5 kJ/mol
    

The result shows that the reaction releases 393.5 kJ/mol of energy, indicating an exothermic reaction.

Let's take another example with a different reaction:

Reaction: N₂(g) + 3H₂(g) → 2NH₃(g)
Given: ΔH°f [N₂(g)] = 0 kJ/mol, ΔH°f [H₂(g)] = 0 kJ/mol, ΔH°f [NH₃(g)] = -45.9 kJ/mol

ΔH = 2(-45.9 kJ/mol) - [0 + 3(0)]
ΔH = -91.8 kJ/mol
    

This calculation shows that 91.8 kJ/mol is released when ammonia is formed from nitrogen and hydrogen, again indicating an exothermic process.

Role of Hess's law

Hess's Law is an important principle in thermochemistry, which states that the total enthalpy change of a chemical reaction is the same regardless of the pathway by which the reaction takes place. This law is helpful in calculating the enthalpy of reactions that are difficult to measure directly. It can be summarized using the following principle:

 ΔH = ∑ΔH (steps in the reaction path) 

Consider the following case of preparation of carbon dioxide by graphite:

C(graphite) + O₂(g) → CO₂(g) ΔH = -393.5 kJ

Now suppose we take a different approach:

C(graphite) + ½O₂(g) → CO(g) ΔH₁ = -110.5 kJ
CO(g) + ½O₂(g) → CO₂(g) ΔH₂ = -283.0 kJ

Applying Hess's Law:
ΔH = ΔH₁ + ΔH₂
ΔH = -110.5 kJ + (-283.0 kJ)
ΔH = -393.5 kJ
    

The result confirms that the enthalpy change remains constant, which demonstrates Hess's law.

Real-world applications of enthalpy and heat of reaction

The concepts of enthalpy and heat of reaction are important in a variety of fields, including:

1. Industrial processes: Controlling energy transformations is critical for efficiency and economy in the manufacturing of chemicals, pharmaceuticals, and materials.
2. Environmental science: Understanding the energy involved in processes such as combustion helps assess their environmental impact.
3. Food science: Calculating the energy emitted or absorbed in food reactions is important for nutrition and conservation.

In addition to shedding light on these applications, enthalpy-related insights also guide research into sustainable energy, such as developing high-efficiency and low-emission fuels.

Conclusion

By studying enthalpy, changes in enthalpy, and heat of reaction, students can better understand how chemical reactions transform matter and energy. This understanding provides a window into both theoretical insights and practical applications, and it forms a foundation for further exploration in chemistry and related sciences.

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