Course Content
Module 1: Fundamental Concepts of Energy
Welcome to the world of chemical energy! Before we can understand how reactions absorb or release heat, we need to establish a common language. In this foundational module, you will be introduced to the key players: the system (the reaction we care about) and the surroundings (everything else). You will learn the critical difference between heat and temperature and be introduced to Enthalpy (H) , the measure of total heat content in a system. By the end of this module, you will understand that in chemistry, we never measure absolute heat—only changes in heat (ΔH) . This concept is the gateway to everything that follows in the course.
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Thermochemistry vs. Thermodynamics
Heat and Temperature
The Concept of Enthalpy (H)
Practice questions
00:20:00
Exothermic and Endothermic Reactions
Have you ever touched a beaker after a reaction and felt it burn your hand, or noticed it turn icy cold? That is thermochemistry in action! This module dives into the two major classes of energy change. You will learn to identify exothermic reactions (ΔH negative) that release heat, like combustion and neutralization, and endothermic reactions (ΔH positive) that absorb heat, like photosynthesis. More importantly, you will learn to draw and interpret energy profile diagrams—a favorite WAEC and JAMB question. These diagrams visually show the energy "hill" that reactions must climb, introducing the concept of Activation Energy (Ea) .
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Exothermic Reactions
Endothermic Reactions
Energy Profile Diagrams
Revision Questions: Endothermic and Exothermic Reactions
00:12:00
Enthalpy Changes of Physical Changes and Chemical Reactions
This is where we get precise. WAEC and JAMB love to test specific definitions. In this module, you will learn the exact language required to score full marks. We will dissect the Standard Enthalpy of Formation (ΔH_f°) —the most important reference point for all energy calculations—and discover why the value for every element in its standard state is zero. You will also master the Enthalpy of Combustion (ΔH_c°) and the constant value associated with the Enthalpy of Neutralization (ΔH_neut) for strong acids and bases. Memorizing these definitions is essential, but understanding them is even more critical.
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Standard Conditions and State Symbols
Principles of Calorimetry
Enthalpy of Formation (ΔH_f°)
Enthalpy of Combustion (ΔH_c°)
Enthalpy of Neutralization (ΔH_neut)
Other Enthalpies
Practice Questions on Enthalpies
00:20:00
Hess’s Law (The Law of Constant Heat Summation)
What if a reaction is too slow, too dangerous, or too difficult to measure in a lab? How do we find its enthalpy change? Enter Hess's Law, one of the most elegant and powerful tools in chemistry. This module teaches you that enthalpy is a "state function"—meaning the path taken from reactants to products doesn't matter; only the start and end points count. You will learn two methods to apply Hess's Law: the algebraic method (manipulating equations) and the energy cycle method (using formation or combustion data). While this is often considered the most challenging topic in thermochemistry, mastering it guarantees you can solve virtually any enthalpy problem thrown at you.
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Statement of Hess’s Law
Applying Hess’s Law (The Algebraic Method)
Energy Cycle Diagrams (The Graphical Method)
Limitations of Bond Enthalpy Calculations
Bond Enthalpies
Chemical reactions involve breaking old bonds and forming new ones. This module looks at energy from the perspective of these bonds. You will learn about mean bond enthalpy—the average energy required to break a specific type of bond. This allows us to estimate the enthalpy change of a reaction using a simple concept: energy must be supplied to break bonds (endothermic), and energy is released when bonds form (exothermic). The formula ΔH = Σ(Bonds Broken) – Σ(Bonds Formed) will become your new best friend. While this method provides an estimate rather than an exact value, it is a quick and useful tool, especially for gaseous reactions. Module 8: Exam-Focused Revision and Problem Solving
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Mean Bond Enthalpy
Calculating ΔH from Bond Enthalpies
Thermochemistry I

Standard Enthalpy of Combustion (\( \Delta H_c^\circ \))

Title: Bomb Calorimeter Used for Measuring Enthalpy of Combustion

Definition

The standard enthalpy of combustion (\( \Delta H_c^\circ \)) is the enthalpy change when one mole of a substance is completely burned in excess oxygen under standard conditions.

Laboratory Experiment to Determine Enthalpy of Combustion

Objective: To determine the enthalpy change when one mole of a liquid fuel (e.g., ethanol) is completely burned. Apparatus:

  • Copper calorimeter (or metal can)
  • Spirit burner containing the liquid fuel
  • Thermometer (\(0.1^\circ\text{C}\) precision)
  • Measuring cylinder
  • Stirrer
  • Balance (accurate to \(0.01\text{ g}\))
  • Clamp and stand

Procedure:

  1. Measure a known volume (e.g., \(100\text{ cm}^3\)) of water into the calorimeter (mass ≈ \(100\text{ g}\)).
  2. Record the initial temperature of the water.
  3. Weigh the spirit burner and record its initial mass.
  4. Place the burner under the calorimeter and light the wick.
  5. Stir continuously; when the temperature rises by \(10\text{–}15^\circ\text{C}\), extinguish the flame.
  6. Record the highest temperature reached and re-weigh the burner.

Assumptions:

  • All heat released is absorbed by the water.
  • Specific heat capacity of water: \(c = 4.18\text{ J g}^{-1}\,^\circ\text{C}^{-1}\).
  • Combustion is complete.

Calculation Steps

1. Temperature change (\( \Delta T \)): \[ \Delta T = \text{Final temperature} – \text{Initial temperature} \] 2. Heat absorbed by water (\( q \)): \[ q = m \cdot c \cdot \Delta T \] 3. Moles of fuel burned (\( n \)): \[ n = \frac{\text{Mass of fuel burned}}{\text{Molar mass of fuel}} \] 4. Enthalpy change per mole: \[ \Delta H_c^\circ = -\frac{q}{n} \] (The negative sign indicates an exothermic reaction)

Worked Example

Problem:

  • Water volume: \(150\text{ cm}^3\)
  • \(\Delta T = 35.5^\circ\text{C} – 22.0^\circ\text{C} = 13.5^\circ\text{C}\)
  • Mass of ethanol burned: \(0.92\text{ g}\)
  • Molar mass of ethanol: \(46.0\text{ g/mol}\)

Solution: Step 1: Heat absorbed (\( q \)): \[ q = 150 \times 4.18 \times 13.5 = 8464.5\text{ J} = 8.4645\text{ kJ} \] Step 2: Moles of ethanol (\( n \)): \[ n = \frac{0.92}{46} = 0.020\text{ mol} \] Step 3: Enthalpy change (\( \Delta H_c^\circ \)): \[ \Delta H_c^\circ = -\frac{8.4645}{0.020} = -423.225\text{ kJ/mol} \] Answer: The enthalpy of combustion of ethanol is approximately \(-423\text{ kJ/mol}\).

Key Characteristics

  • Always negative: Combustion reactions are exothermic.
  • Complete combustion: Carbon \( \rightarrow CO_2(g) \), Hydrogen \( \rightarrow H_2O(l) \).

Examples

Methane

Formula: \(CH_4(g)\)

Combustion Equation:

\(CH_4(g) + 2O_2(g) \rightarrow CO_2(g) + 2H_2O(l)\)

\( \Delta H_c^\circ \): -890 kJ/mol

Ethanol

Formula: \(C_2H_5OH(l)\)

Combustion Equation:

\(C_2H_5OH(l) + 3O_2(g) \rightarrow 2CO_2(g) + 3H_2O(l)\)

\( \Delta H_c^\circ \): -1367 kJ/mol

Hydrogen

Formula: \(H_2(g)\)

Combustion Equation:

\(H_2(g) + \frac{1}{2}O_2(g) \rightarrow H_2O(l)\)

\( \Delta H_c^\circ \): -286 kJ/mol

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