This study guide covers fundamental thermochemistry concepts including definitions of temperature, heat, specific heat, and various enthalpy changes. It details exothermic and endothermic reactions, thermochemical equations, and specific enthalpy types like molar enthalpy of formation and combustion. Students should be prepared to apply formulas, interpret reaction pathways, and understand Hess's Law for calculating overall enthalpy changes. Key distinctions include exothermic versus endothermic processes and the stability of reactants versus products.
Basic Thermochemistry Definitions
This section defines fundamental concepts in thermochemistry, including heat and enthalpy.
Concept
Temperature
Measure of average kinetic energy of the particles in a sample of matter.
Concept
Heat
Energy transferred between samples of matter because of a difference in temperature.
Concept
Specific Heat (cp)
Amount of energy required to raise the temperature of 1g of substance by 1°C (J/(g*k)).
Formula
q = c_p * m * ΔT
Energy lost or gained (q) equals specific heat (cp) times mass (m) times change in temperature (ΔT).
Concept
Enthalpy Change (ΔH)
Amount of energy absorbed by a system as heat during a process at constant pressure.
Concept
Enthalpy of Reaction
Quantity of energy transferred as heat during a chemical reaction.
Concept
Exothermic Reaction
A reaction that releases heat to the surroundings.
Exothermic Reaction Pathway
This section illustrates the energy profile for an exothermic reaction.
- 1
Initial State
Reactants begin at a higher initial enthalpy.
- 2
Energy Evolved
As the reaction proceeds, energy is released as heat to the surroundings.
- 3
Final State
Products end at a lower final enthalpy.
Key Characteristics
For an exothermic reaction, the enthalpy change (ΔH) is negative, indicating a net release of energy. The products are in a lower energy state and are therefore more stable than the reactants.
Endothermic Reaction and Pathway
This section defines endothermic reactions and illustrates their energy profile.
Concept
Endothermic Reaction
A reaction that absorbs heat from the surroundings.
- 1
Initial State
Reactants begin at a lower initial enthalpy.
- 2
Energy Absorbed
As the reaction proceeds, energy is absorbed as heat from the surroundings.
- 3
Final State
Products end at a higher final enthalpy.
| Endothermic Reaction | Exothermic Reaction | |
|---|---|---|
| Heat Flow | Absorbs heat from surroundings | Releases heat to surroundings |
| ΔH Sign | ΔH is positive | ΔH is negative |
| Stability | Reactants are stable in a low energy state | Products are stable in a low energy state |
| Enthalpy Change | Final enthalpy > Initial enthalpy | Final enthalpy < Initial enthalpy |
Thermochemical Equation
This section defines thermochemical equations and outlines rules for their use.
Concept
Thermochemical Equation
A balanced chemical equation that includes the amount of energy released or absorbed as heat (ΔH).
Examples:
- 2H2 (g) + O2 (g) → 2H2O (g), ΔH = -483.6kJ
- H2 (g) + 1/2 O2 (g) ⇒ H2O (g), ΔH = -241.8kJ
| Coefficients represent | Coefficients do NOT represent | |
|---|---|---|
| Quantity | Numbers of moles of reactants and products | Numbers of molecules |
- Coefficients can be fractions when necessary.
- The physical state (g, l, s, aq) of each reactant and product must be included.
- The enthalpy change (ΔH) is directly proportional to the number of moles of substances undergoing change. (e.g., decomposing 2 mol of water needs twice the enthalpy of 1 mol).
- The value of ΔH is usually not significantly influenced by changing temperature.
Molar Enthalpy of Formation (ΔH_f)
This section defines molar enthalpy of formation and provides examples.
Concept
Molar Enthalpy of Formation (ΔH_f)
The enthalpy change that occurs when one mole of a compound is formed from its elements in their standard state (25 Celsius degree, 1 atm).
Examples illustrating exothermic and endothermic formation reactions:
Concept
Exothermic Formation Example
C(s) + O2(g) ⇒ CO2(g)
ΔH_f = -393.5kJ (heat is released).
Concept
Endothermic Formation Example
1/2 N2(g) + 1/2 O2(g) ⇒ NO(g)
ΔH_f = +90.29kJ (heat is absorbed).
Enthalpy of Combustion (ΔH_c)
This section defines enthalpy of combustion and provides examples.
Concept
Enthalpy of Combustion (ΔH_c)
Enthalpy change that occurs during the combustion of one mole of a substance in the presence of oxygen.
Examples of combustion reactions:
- CO (g) + 1/2 O2(g) → CO2(g), ΔH_c = -283.0kJ
- C4H10 (g) + 13/2 O2(g) ⇒ 4CO2(g) + 5H2O(l), ΔH_c = -2877.6kJ
Hess's Law and Calculations
This section introduces Hess's Law and demonstrates its application in calculations.
Concept
Hess's Law
The overall enthalpy change in a reaction is equal to the sum of the enthalpy changes for the individual steps in the process, regardless of the pathway taken.
Calculating Enthalpy of Formation for Methane Gas (CH₄)
- 1
Target Reaction
C(s) + 2H2(g) → CH₄(g) ΔH = ?
- 2
Given Reactions
1. C(s) + O2(g) → CO2(g) ΔH = −393.5 kJ
2. H2(g) + 1/2 O2(g) → H2O(l) ΔH = −285.8 kJ
3. CH₄(g) + 2O2(g) → CO2(g) + 2H2O(l) ΔH = −890.8 kJ
- 3
Step 1: Manipulate Reaction 1
Keep as is to get C(s) on the reactant side:
C(s) + O2(g) ⇒ CO2(g) ΔH = -393.5kJ
- 4
Step 2: Manipulate Reaction 2
Multiply by 2 to get 2H2(g) on the reactant side and 2H2O(l) on the product side:
2H2(g) + O2(g) → 2H2O(l) ΔH = *2 (-285.8 kJ) = -571.6kJ**
- 5
Step 3: Manipulate Reaction 3
Reverse the reaction to get CH₄(g) on the product side. This changes the sign of ΔH:
CO2(g) + 2H2O(l) ⇒ CH4(g) + 2O2(g) ΔH = +890.8kJ
- 6
Step 4: Sum the Manipulated Reactions
Add the modified reactions and their ΔH values, canceling common species on opposite sides:
C(s) + O2(g) + 2H2(g) + O2(g) + CO2(g) + 2H2O(l) ⇒ CO2(g) + 2H2O(l) + CH4(g) + 2O2(g)
After canceling: C(s) + 2H2(g) ⇒ CH4(g)
Answer
ΔH = (-393.5 kJ) + (-571.6 kJ) + (+890.8 kJ) = -74.3 kJ