Class 11 Chemistry – Chapter 6: Energetics (FBISE)
This section provides complete, exam-oriented notes for Class 11 Chemistry Chapter 6 – Energetics strictly according to the Federal Board (FBISE) syllabus. Concepts are explained in a clear and student-friendly manner for easy understanding and exam preparation.
Key topics include Thermochemistry, Standard Conditions, Standard Heats, Born-Haber Cycle, Hess's Law, Lattice Energy, Hydration Energy, and the factors affecting entropy. Detailed examples, calculations, and step-by-step solutions are provided for effective learning.
Students can also access video lectures, solved numericals, MCQs, test series, and live classes for this chapter on our official YouTube channel and stay updated through our WhatsApp channel.
6.1 Energy in Chemical Reactions
Chemical reactions involve the breaking of old bonds and the formation of new bonds. This process leads to the evolution (release) or absorption of heat energy.
- Bond Breaking: Always consumes energy (Endothermic).
- Bond Making: Always releases energy (Exothermic).
The net energy change depends on the balance between these two processes. If bond making releases more energy than bond breaking consumes, energy is transferred to the surroundings.
6.1.1 Units of Thermal Energy
Thermal energy is the energy in a system due to the movement of its atoms and molecules.
- Joule (J): The SI unit of heat or thermal energy.
- Calorie (cal): The amount of heat required to raise the temperature of 1 gram of water from $14.5^\circ\text{C}$ to $15.5^\circ\text{C}$.
Conversion: $1 \text{ Calorie} = 4.18 \text{ Joules}$
6.1.2 Thermochemical Reactions
Thermochemistry is the branch of chemistry dealing with heat changes in chemical reactions. A Thermochemical Equation is a balanced chemical equation that includes the heat change ($\Delta H^\circ$).
1. Exothermic Reactions
Reactions that proceed with the evolution (release) of heat. Energy is transferred from the system to the surroundings. $\Delta H$ is represented as a negative value.
- Example: $C_{(s)} + O_{2(g)} \rightarrow CO_{2(g)} \quad \Delta H^\circ = -393.5 \text{ kJ}$
2. Endothermic Reactions
Reactions that proceed with the absorption of heat. Energy is transferred from the surroundings to the system. $\Delta H$ is represented as a positive value.
- Example: $N_{2(g)} + O_{2(g)} \rightarrow 2NO_{(g)} \quad \Delta H^\circ = +180.5 \text{ kJ}$
Q&A and Concept Assessment
Relevant Questions & Answers
Q1: Why is energy either absorbed or evolved during a chemical reaction?A: Because chemical reactions involve breaking existing bonds (which requires energy) and forming new bonds (which releases energy). The difference between these two determines if the net energy is absorbed or evolved.
Q2: Define a Calorie.A: A calorie is the thermal energy required to raise the temperature of one gram of water by $1^\circ\text{C}$ (specifically from $14.5^\circ\text{C}$ to $15.5^\circ\text{C}$).
Concept Assessment Exercise 6.1 - Solutions
Classify the following processes as exothermic or endothermic:
| Process | Classification | Reasoning |
|---|---|---|
| (a) Freezing of water | Exothermic | Heat is released as liquid turns to solid. |
| (b) Combustion of methane | Exothermic | Burning fuels always releases heat energy. |
| (c) Sublimation of dry ice | Endothermic | Energy is absorbed to turn solid $CO_2$ directly into gas. |
| (d) $H_2O_{(g)} \rightarrow H_2O_{(l)}$ | Exothermic | Condensation releases heat as molecules slow down. |
| (e) Decomposition of limestone | Endothermic | Heat must be supplied to break down $CaCO_3$. |
6.1.3 Heat of Reaction
The energy change in a reaction is directly proportional to the amount of substances that react. When reported in terms of molar quantities as shown in a balanced chemical equation, this energy is called the heat of reaction.
- Standard Enthalpy Change ($\Delta H^\circ$): This is the heat of reaction measured at $25^\circ\text{C}$ ($298\text{K}$) and $1 \text{ atm}$ pressure.
- Reversibility: If a reaction is exothermic in one direction, it is endothermic in the reverse direction. While the magnitude of $\Delta H^\circ$ stays the same, the sign changes.
Examples of Reversible Thermochemical Equations:
- Forward: $C_{(s)} + O_{2(g)} \rightarrow CO_{2(g)} \quad \Delta H^\circ = -393.5\text{kJ}$
- Reverse: $CO_{2(g)} \rightarrow C_{(s)} + O_{2(g)} \quad \Delta H^\circ = +393.5\text{kJ}$
6.1.4 Relation between Enthalpy Change and Heat of Reaction
Since most reactions occur at constant pressure, heat change is equated to the change in enthalpy ($\Delta H$). It is defined as the difference between the enthalpies of the products and the reactants:
$$\Delta H = H_{\text{products}} - H_{\text{reactants}}$$
- Exothermic: $\Delta H$ is negative because heat is released to the surroundings.
- Endothermic: $\Delta H$ is positive because heat is absorbed from the surroundings.
6.1.5 & 6.1.6 Standard States and Conditions
Enthalpy values vary with conditions, so standardized values are calculated when substances are in their standard state.
Standard State Conditions:
- Gases: Pressure of $1 \text{ atm}$.
- Elements/Compounds: Most stable physical state at $1 \text{ atm}$ and $25^\circ\text{C}$ ($298\text{K}$).
- Aqueous Solutions: Concentration of $1\text{M}$.
Relevant Questions & Answers
Q1: What happens to the sign and magnitude of enthalpy when a chemical reaction is reversed?A: The magnitude of the enthalpy ($\Delta H^\circ$) remains the same, but the sign changes (e.g., negative becomes positive).
Q2: Define "Standard Enthalpy Change".A: It is the heat of reaction measured under standard conditions: $25^\circ\text{C}$ ($298\text{K}$) and $1 \text{ atmospheric pressure}$.
Q3: Explain why $\Delta H$ is negative for combustion reactions.A: Combustion is an exothermic process where heat is released by the system to the surroundings, meaning the products have lower enthalpy than the reactants.
Q4: Interpret the following equation: $H_{2(g)} + I_{2(g)} \rightarrow 2HI_{(g)} \quad \Delta H^\circ = +53.8\text{kJ}$A: This equation shows that when $1 \text{ mole}$ of hydrogen gas reacts with $1 \text{ mole}$ of iodine gas to produce $2 \text{ moles}$ of hydrogen iodide at $25^\circ\text{C}$ and $1 \text{ atm}$, $53.8\text{kJ}$ of heat is absorbed.
6.1.3 Heat of Reaction
The amount of energy released or absorbed in a chemical reaction is proportional to the quantity of the reacting substances.
- Definition: The heat evolved or absorbed when molar quantities of reactants and products are exactly as shown in the balanced chemical equation is called the heat of reaction.
- Standard Enthalpy Change ($\Delta H^\circ$): This refers to the heat of reaction measured at $25^\circ\text{C}$ ($298\text{K}$) and $1 \text{ atmospheric pressure}$.
- Reversibility: When a reaction is reversed, the magnitude of $\Delta H^\circ$ remains identical, but the algebraic sign changes (e.g., an exothermic forward reaction becomes endothermic in reverse).
6.1.4 Enthalpy Change ($\Delta H$)
At constant pressure, the heat change of a reaction is defined as the difference between the enthalpies of the products and reactants.
$$\Delta H = H_{\text{products}} - H_{\text{reactants}}$$
- Exothermic: $\Delta H$ is negative because heat is released to the surroundings.
- Endothermic: $\Delta H$ is positive because heat is absorbed from the surroundings.
6.1.6 Conditions for Standard States
Standardized enthalpy values are calculated when substances are in their standard states:
- Gases: Pressure is $1 \text{ atm}$.
- Elements/Compounds: The most stable physical state at $1 \text{ atm}$ and $25^\circ\text{C}$ ($298\text{K}$).
- Aqueous Solutions: Concentration of $1\text{M}$.
Types of Standard Enthalpies
| Type | Symbol | Definition |
|---|---|---|
| Standard Enthalpy of Formation | $\Delta H_f^\circ$ | Enthalpy change accompanying the formation of one mole of a compound from its constituent elements in their standard states. |
| Standard Enthalpy of Combustion | $\Delta H_c^\circ$ | Enthalpy change when one mole of a substance is completely burnt in excess oxygen. |
| Standard Enthalpy of Atomization | $\Delta H_{at}^\circ$ | Enthalpy change when one mole of gaseous atoms are formed from its element. |
Relevant Questions & Answers
Q1: Interpret the equation $C_{(s)} + O_{2(g)} \rightarrow CO_{2(g)} \quad \Delta H^\circ = -393.5 \text{ kJ}$ in terms of molar quantities.A: This equation indicates that $1 \text{ mole}$ of solid carbon ($12\text{g}$) reacts with $1 \text{ mole}$ of oxygen gas ($32\text{g}$) to produce $1 \text{ mole}$ of $CO_2$ gas ($44\text{g}$), evolving $393.5 \text{ kJ}$ of heat.
Q2: If the formation of $HI$ gas is endothermic ($+53.8 \text{ kJ}$), what is the enthalpy for its decomposition?A: The decomposition (reverse reaction) would be exothermic with an enthalpy change of $-53.8 \text{ kJ}$.
Q3: Define the Standard Enthalpy of Atomization with an example.A: It is the enthalpy change when one mole of gaseous atoms is formed from its element.Example: $\frac{1}{2}H_{2(g)} \rightarrow H_{(g)} \quad \Delta H_{at}^\circ = +218 \text{ kJ mole}^{-1}$.
Q4: Why is the enthalpy of combustion ($\Delta H_c^\circ$) always negative?A: Combustion is an exothermic process because it always involves the release of heat to the surroundings.
6.1.3 Heat of Reaction
The amount of energy released or absorbed in a chemical reaction depends on the quantities of substances that react.
- Definition: Heat of reaction is the amount of heat evolved or absorbed when molar quantities of reactants and products are exactly as shown in a balanced chemical equation.
- Standard Enthalpy Change ($\Delta H^\circ$): This is the heat of reaction measured at $25^\circ\text{C}$ ($298\text{K}$) and $1 \text{ atmospheric pressure}$.
- Sign Convention: If a reaction is exothermic in one direction, it is endothermic in the reverse direction. When reversed, the magnitude of $\Delta H^\circ$ stays the same, but the sign changes.
6.1.4 Relation between Enthalpy Change and Heat of Reaction
For reactions occurring at constant pressure, heat change is equated to the change in enthalpy ($\Delta H$).
$$\Delta H = H_{\text{products}} - H_{\text{reactants}}$$
- Exothermic: $\Delta H$ is negative because heat is released by the system.
- Endothermic: $\Delta H$ is positive because heat is absorbed by the system.
6.1.6 Conditions for Standard Heat of Reaction
- Standard state for a gas is $1 \text{ atm}$.
- Standard state for an element or compound is the most stable physical state at $1 \text{ atm}$ and $25^\circ\text{C}$.
- Standard state for a substance in aqueous solution is $1\text{M}$ concentration.
Types of Standard Enthalpies
| Enthalpy Type | Definition | Example Equation |
|---|---|---|
| Formation ($\Delta H_f^\circ$) | The change that accompanies the formation of one mole of a compound from its elements in their standard states. | $H_{2(g)} + \frac{1}{2}O_{2(g)} \rightarrow H_2O_{(l)} \quad \Delta H_f^\circ = -285.8 \text{ kJ mole}^{-1}$ |
| Combustion ($\Delta H_c^\circ$) | The change when one mole of a substance is completely burnt in excess oxygen. | $CH_{4(g)} + 2O_{2(g)} \rightarrow CO_{2(g)} + 2H_2O_{(l)} \quad \Delta H_c^\circ = -890.4 \text{ kJ mole}^{-1}$ |
| Atomization ($\Delta H_{at}^\circ$) | The change when one mole of gaseous atoms are formed from its element. | $\frac{1}{2}H_{2(g)} \rightarrow H_{(g)} \quad \Delta H_{at}^\circ = +218 \text{ kJ mole}^{-1}$ |
| Neutralization ($\Delta H_n^\circ$) | Heat evolved when one mole of $H^+$ ions from an acid combine with one mole of $OH^-$ ions from a base to form water. | $H^+_{(aq)} + OH^-_{(aq)} \rightarrow H_2O_{(l)} \quad \Delta H_n^\circ = -57.4 \text{ kJ mole}^{-1}$ |
| Solution ($\Delta H_{sol}^\circ$) | The change when one mole of a substance is dissolved in so much solvent that further dilution results in no detectable heat change. | $NH_4Cl_{(s)} \xrightarrow{H_2O} NH^+_{4(aq)} + Cl^-_{(aq)} \quad \Delta H_{sol}^\circ = +15.1 \text{ kJ mole}^{-1}$ |
6.2 BOND ENERGY
Chemical reactions involve breaking old bonds (requires energy) and forming new bonds (releases energy).
- Bond Dissociation Energy: Energy required to break one mole of a particular bond to form neutral atoms.
- Calculation: The enthalpy change of a reaction is the sum of bond dissociation energies of reactants minus the sum of bond energies of products.
$$\Delta H = \sum H_{\text{bonds broken}} - \sum H_{\text{bonds formed}}$$
Relevant Questions & Answers
Q1: What is Lattice Energy?A: Lattice energy is the energy released when one mole of an ionic solid compound is formed from its constituent gaseous ions (e.g., $Na^+_{(g)} + Cl^-_{(g)} \rightarrow NaCl_{(s)} \quad \Delta H = -787 \text{ kJ mole}^{-1}$).
Q2: Why is the enthalpy of neutralization for strong acids and bases usually constant?A: Strong acids and bases ionize completely in aqueous solution. The only significant change is the formation of water from $H^+$ and $OH^-$ ions, which always evolves $57.4 \text{ kJ}$ per mole of water formed.
Q3: How is average bond energy calculated for a molecule like methane ($CH_4$)?A: By taking the total energy required to break all bonds in the molecule and dividing by the number of bonds. For $CH_4$, breaking $4$ moles of C-H bonds requires $1662 \text{ kJ}$, so the average C-H bond energy is $\frac{1662}{4} = +415.5 \text{ kJ/mole}$.
Q4: Define Standard Enthalpy of first Electron Affinity.A: It is the energy released when one mole of gaseous atoms gains an electron to form one mole of gaseous ions with a single negative charge (e.g., $Cl_{(g)} + 1e^- \rightarrow Cl^-_{(g)} \quad \Delta H_{EA}^\circ = -349 \text{ kJ/mole}$).
6.1.3 Heat of Reaction
The amount of energy released or absorbed in a chemical reaction depends on the amounts of substances that react.
- Definition: The amount of heat evolved or absorbed in a chemical reaction, when molar quantities of reactants and products are the same as shown in a balanced chemical equation, is called the heat of reaction.
- Standard Enthalpy Change ($\Delta H^\circ$): Heat of reaction measured at $25^\circ\text{C}$ ($298\text{K}$) and $1 \text{ atmospheric pressure}$.
- Reversibility: If a reaction is exothermic in one direction, it is endothermic in the reverse direction. When a reaction is reversed, the magnitude of $\Delta H^\circ$ remains the same but the sign changes.
- Example: $C_{(s)} + O_{2(g)} \rightarrow CO_{2(g)} \quad \Delta H^\circ = -393.5\text{kJ}$
- Reverse: $CO_{2(g)} \rightarrow C_{(s)} + O_{2(g)} \quad \Delta H^\circ = +393.5\text{kJ}$
6.1.4 Relation between Enthalpy Change and Heat of Reaction
Since most reactions occur at constant pressure, heat change is equated to the change in enthalpy ($\Delta H$). It is defined as the difference between the enthalpies of the products (final state) and reactants (initial state):
$$\Delta H = H_{\text{products}} - H_{\text{reactants}}$$
- Exothermic process: $\Delta H$ is negative because heat is released by the system.
- Endothermic process: $\Delta H$ is positive because heat is absorbed by the system.
6.1.6 Conditions for Standard Heat of Reaction
- Standard state for a gas is $1 \text{ atm}$.
- Standard state for an element or compound is the most stable physical state at $1 \text{ atm}$ and $25^\circ\text{C}$ ($298\text{K}$).
- Standard state for a substance in aqueous solution is $1\text{M}$ concentration.
Types of Standard Enthalpies
| Enthalpy Type | Definition | Example |
|---|---|---|
| Formation ($\Delta H_f^\circ$) | Enthalpy change accompanying the formation of one mole of a compound from its elements in their standard states. | $H_{2(g)} + \frac{1}{2}O_{2(g)} \rightarrow H_2O_{(l)} \quad \Delta H_f^\circ = -285.8 \text{ kJ mole}^{-1}$ |
| Combustion ($\Delta H_c^\circ$) | Enthalpy change when one mole of a substance is completely burnt in excess oxygen. | $CH_{4(g)} + 2O_{2(g)} \rightarrow CO_{2(g)} + 2H_2O_{(l)} \quad \Delta H_c^\circ = -890.4 \text{ kJ mole}^{-1}$ |
| Atomization ($\Delta H_{at}^\circ$) | Enthalpy change when one mole of gaseous atoms are formed from its element. | $\frac{1}{2}H_{2(g)} \rightarrow H_{(g)} \quad \Delta H_{at}^\circ = +218 \text{ kJ mole}^{-1}$ |
| Neutralization ($\Delta H_n^\circ$) | Heat evolved when one mole of $H^+$ ions from an acid combine with one mole of $OH^-$ ions from a base to form water. | $H^+_{(aq)} + OH^-_{(aq)} \rightarrow H_2O_{(l)} \quad \Delta H_n^\circ = -57.4 \text{ kJ mole}^{-1}$ |
| Solution ($\Delta H_{sol}^\circ$) | Enthalpy change when one mole of a substance is dissolved in so much solvent that further dilution results in no detectable heat change. | $NH_4Cl_{(s)} \xrightarrow{H_2O} NH^+_{4(aq)} + Cl^-_{(aq)} \quad \Delta H_{sol}^\circ = +15.1 \text{ kJ mole}^{-1}$ |
6.2 BOND ENERGY
Breaking bonds requires energy, and forming bonds releases energy.
- Bond Dissociation Energy: The amount of energy required to break one mole of a particular bond to form neutral atoms.
- Calculation: $\Delta H = \sum H_{\text{(bonds broken)}} - \sum H_{\text{(bonds formed)}}$.
6.4 HESS'S LAW
Enthalpy is a state function; the enthalpy change in a chemical reaction is independent of the path followed.
Definition: The enthalpy change in a chemical reaction is the same whether the reaction takes place in a single step or in several steps.
$$\sum \Delta H \text{ (Cycle)} = 0$$
Relevant Questions & Answers
Q1: Interpret the following thermochemical equation: $H_{2(g)} + I_{2(g)} \rightarrow 2HI_{(g)} \quad \Delta H^\circ = +53.8\text{kJ}$A: This equation shows that at $25^\circ\text{C}$ and $1 \text{ atmospheric pressure}$, the reaction between gaseous hydrogen and iodine to form $2 \text{ moles}$ of $HI$ gas is endothermic, absorbing $53.8\text{kJ}$ of heat.
Q2: Define Lattice Energy and provide an example.
A: Lattice energy is the energy released when one mole of an ionic solid compound is formed from its constituent gaseous ions.
Example: $Na^+_{(g)} + Cl^-_{(g)} \rightarrow NaCl_{(s)} \quad \Delta H = -787 \text{ kJ mole}^{-1}$.
Q3: Why is the experimental bond energy of $HCl$ ($431.5 \text{ kJ/mol}$) different from the theoretical value ($339 \text{ kJ/mol}$)?
A: Theoretical values are averages and do not account for specific molecular environments, such as unequal distribution of electron pairs due to differences in electronegativity values.
Q4: What is the Standard Enthalpy of first Electron Affinity?
A: It is the energy released when one mole of gaseous atoms gains an electron to form one mole of gaseous ions with a single negative charge.
Example: $Cl_{(g)} + 1e^- \rightarrow Cl^-_{(g)} \quad \Delta H_{EA}^\circ = -349 \text{ kJ/mole}$.
Q5: According to Hess's Law, if a reaction $A \rightarrow B$ happens in one step ($\Delta H$) or via $A \rightarrow C \rightarrow D \rightarrow B$ ($\Delta H_1, \Delta H_2, \Delta H_3$), what is the relationship between these values?
A: According to Hess's Law, $\Delta H = \Delta H_1 + \Delta H_2 + \Delta H_3$.
6.1 Energy in Chemical Reactions
Chemical reactions involve the breaking of old bonds and the formation of new bonds. This process is accompanied by the absorption or evolution of energy in the form of heat.
- Bond Breaking: Always consumes energy.
- Bond Formation: Always releases energy.
- Net Transfer: Energy is transferred to the surroundings if formation releases more energy than breaking consumes. Conversely, energy is absorbed from surroundings if breaking consumes more energy.
6.1.1 Units of Thermal Energy
Thermal energy is the energy in a system due to the movement of its atoms and molecules.
- Joule (J): The SI unit for heat or thermal energy.
- Calorie (cal): The amount of thermal energy required to raise the temperature of $1 \text{ gram}$ of water from $14.5^\circ\text{C}$ to $15.5^\circ\text{C}$.
- Conversion: $1 \text{ Calorie} = 4.18 \text{ Joules}$.
6.1.2 Thermochemical Reactions
A chemical reaction that proceeds with the evolution or absorption of heat is a thermochemical reaction. The branch of chemistry dealing with these changes is Thermochemistry.
- Exothermic Reactions: Heat is evolved; the system transfers energy to the surroundings. $\Delta H$ is negative. Example: Burning fuels.
- Endothermic Reactions: Heat is absorbed; energy is transferred from surroundings to the system. $\Delta H$ is positive. Example: Decomposition of limestone.
6.1.4 Enthalpy Change ($\Delta H$)
Enthalpy change is defined as the difference between the enthalpies of products and reactants.
$$\Delta H = H_{\text{products}} - H_{\text{reactants}}$$
6.1.6 Standard Conditions for Heat of Reaction
- Gases: Pressure of $1 \text{ atm}$.
- Elements/Compounds: Most stable physical state at $1 \text{ atm}$ and $25^\circ\text{C}$ ($298 \text{ K}$).
- Aqueous Solutions: Concentration of $1 \text{ M}$.
Types of Standard Enthalpies
| Type | Symbol | Definition |
|---|---|---|
| Formation | $\Delta H_f^\circ$ | Enthalpy change when one mole of a compound is formed from its elements in their standard states. |
| Combustion | $\Delta H_c^\circ$ | Enthalpy change when one mole of a substance is completely burnt in excess oxygen. |
| Atomization | $\Delta H_{at}^\circ$ | Enthalpy change when one mole of gaseous atoms are formed from its element. |
| Neutralization | $\Delta H_n^\circ$ | Heat evolved when one mole of $H^+$ ions from an acid combine with one mole of $OH^-$ ions from a base to form water. |
| Solution | $\Delta H_{sol}^\circ$ | Enthalpy change when one mole of a substance is dissolved in enough solvent that further dilution causes no detectable heat change. |
| Lattice Energy | $\Delta H$ | Energy released when one mole of an ionic solid is formed from its constituent gaseous ions. |
6.2 Bond Energy & 6.4 Hess's Law
- Bond Dissociation Energy: Energy required to break one mole of a particular bond to form neutral atoms.
- Reaction Enthalpy Calculation: $$\Delta H = \sum H_{\text{bonds broken}} - \sum H_{\text{bonds formed}}$$
- Hess's Law: Enthalpy change in a chemical reaction is independent of the path followed; it is the same whether the reaction occurs in one step or several.
Relevant Questions & Answers
Q1: Why does a reaction's $\Delta H^\circ$ sign change when it is reversed?A: Because if a reaction is exothermic (releasing heat) in one direction, it must absorb the same amount of heat to return to its original state (endothermic), keeping the magnitude identical but flipping the sign.
Q2: Calculate the average C-H bond energy in methane ($CH_4$) if the total energy absorbed to break 4 moles of bonds is $+1662 \text{ kJ}$.A: The average bond energy is the total energy divided by the number of bonds: $\frac{+1662 \text{ kJ}}{4 \text{ moles}} = +415.5 \text{ kJ/mole}$.
Q3: How do you classify the "Freezing of water" based on energy?A: Freezing is an exothermic process because the water system must release heat to the surroundings to transition from liquid to solid.
Q4: State the formula used to calculate the Enthalpy of Reaction from Enthalpies of Formation.A: $$\Delta H^\circ = \sum \text{coeff} \Delta H_f^\circ (\text{products}) - \sum \text{coeff} \Delta H_f^\circ (\text{reactants})$$
6.1 Energy in Chemical Reactions
Chemical reactions involve the breaking of old bonds and the formation of new bonds.
- Bond Breaking: Always consumes energy.
- Bond Making: Always releases energy.
- Net Transfer: Energy is transferred from the system to the surroundings if the energy released by bond formation is greater than that consumed by bond breaking. Conversely, if bond breaking consumes more energy, it is transferred from the surroundings to the system.
6.1.1 Units of Thermal Energy
- Joule (J): The SI unit of heat or thermal energy.
- Calorie: Defined as the heat required to raise the temperature of $1 \text{ gram}$ of water from $14.5^\circ\text{C}$ to $15.5^\circ\text{C}$.
- Conversion: $1 \text{ Calorie} = 4.18 \text{ Joules}$.
6.1.2 Thermochemical Reactions
A reaction that proceeds with the evolution or absorption of heat is a thermochemical reaction.
- Exothermic Reactions: Heat is evolved (released). $\Delta H$ is negative. Examples include the burning of fuels.
- Endothermic Reactions: Heat is absorbed from the surroundings. $\Delta H$ is positive. Examples include the decomposition of limestone.
6.1.3 & 6.1.4 Heat of Reaction and Enthalpy Change
The amount of heat evolved or absorbed when molar quantities of reactants react as shown in a balanced equation is called the heat of reaction.
- Standard Enthalpy Change ($\Delta H^\circ$): Heat of reaction measured at $25^\circ\text{C}$ ($298\text{K}$) and $1 \text{ atmospheric pressure}$.
- Enthalpy Change ($\Delta H$): Defined as the difference between the enthalpies of products and reactants.
$$\Delta H = H_{\text{products}} - H_{\text{reactants}}$$
6.1.6 Conditions for Standard States
- Gases: Pressure of $1 \text{ atm}$.
- Elements/Compounds: Most stable physical state at $1 \text{ atm}$ and $25^\circ\text{C}$.
- Aqueous Solutions: Concentration of $1\text{M}$.
Types of Standard Enthalpies
| Type | Symbol | Definition |
|---|---|---|
| Formation | $\Delta H_f^\circ$ | Enthalpy change when one mole of a compound is formed from its elements in their standard states. |
| Combustion | $\Delta H_c^\circ$ | Enthalpy change when one mole of a substance is completely burnt in excess oxygen. |
| Atomization | $\Delta H_{at}^\circ$ | Enthalpy change when one mole of gaseous atoms are formed from its element. |
| Neutralization | $\Delta H_n^\circ$ | Heat evolved when one mole of $H^+$ ions from an acid combine with one mole of $OH^-$ ions from a base to form water. |
| Solution | $\Delta H_{sol}^\circ$ | Enthalpy change when one mole of a substance is dissolved in enough solvent that further dilution results in no heat change. |
| Lattice Energy | $\Delta H_l^\circ$ | Energy released when one mole of an ionic solid is formed from its constituent gaseous ions. |
6.2 Bond Energy & 6.4 Hess's Law
- Bond Dissociation Energy: Energy required to break one mole of a particular bond to form neutral atoms.
- Hess's Law: The enthalpy change in a chemical reaction is the same whether it occurs in one step or several.
$$\sum \Delta H \text{ (Cycle)} = 0$$
Relevant Questions & Answers
Q1: Why is $\Delta H$ negative for exothermic reactions?A: In exothermic reactions, heat is released by the system to the surroundings, meaning the enthalpy of the products is lower than the enthalpy of the reactants.
Q2: If a forward reaction has $\Delta H^\circ = -393.5 \text{ kJ}$, what is $\Delta H^\circ$ for the reverse reaction?
A: The magnitude remains the same, but the sign changes. Thus, $\Delta H^\circ = +393.5 \text{ kJ}$.
Q3: Define the Standard Enthalpy of first Electron Affinity.
A: It is the energy released when one mole of gaseous atoms gains an electron to form one mole of gaseous ions with a single negative charge.
Q4: How is the average bond energy of a C-H bond in methane ($CH_4$) calculated?
A: By dividing the total energy required to break all four C-H bonds ($1662 \text{ kJ}$) by four: $\frac{1662}{4} = 415.5 \text{ kJ/mole}$.
Q5: Classify the "Sublimation of dry ice" as exothermic or endothermic.
A: It is endothermic because heat must be absorbed from the surroundings for solid $CO_2$ to turn into gas.
6.1 Energy in Chemical Reactions
Chemical reactions involve the breaking of old bonds and the formation of new bonds. This process leads to the absorption or release of energy in the form of heat.
- Bond Breaking: Always consumes energy.
- Bond Making: Always releases energy.
- Energy Transfer: If energy released by bond formation is greater than energy consumed by bond breaking, energy is transferred from the system to the surroundings. Conversely, if bond breaking consumes more energy, heat is transferred from the surroundings to the system.
6.1.1 Units of Thermal Energy
Thermal energy is the energy in an object or system due to the movement of its molecules and atoms.
- Joule (J): The unit of heat or thermal energy used in the SI system.
- Calorie (cal): The amount of heat required to raise the temperature of $1 \text{ gram}$ of water from $14.5^\circ\text{C}$ to $15.5^\circ\text{C}$.
- Conversion: $1 \text{ Calorie} = 4.18 \text{ Joules}$.
6.1.2 Thermochemical Reactions
A chemical reaction which proceeds with the evolution or absorption of heat is called a thermochemical reaction. A balanced chemical equation showing this heat change is a thermochemical equation.
Shutterstock1. Exothermic Reactions
Reactions that proceed with the evolution (release) of heat. The system transfers energy to the surroundings.
- Example: $C_{(s)} + O_{2(g)} \rightarrow CO_{2(g)} \quad \Delta H^\circ = -393.5 \text{ kJ}$.
2. Endothermic Reactions
Reactions that proceed with the absorption of heat. Heat is transferred from the surroundings to the system.
- Example: $N_{2(g)} + O_{2(g)} \rightarrow 2NO_{(g)} \quad \Delta H^\circ = +180.5 \text{ kJ}$.
6.1.3 Heat of Reaction & Enthalpy
The energy change in a reaction is proportional to the amounts of substances that react. When reported for molar quantities as shown in a balanced equation, it is called the heat of reaction.
- Standard Enthalpy Change ($\Delta H^\circ$): Heat of reaction measured at $25^\circ\text{C}$ ($298 \text{ K}$) and $1 \text{ atm}$ pressure.
- Enthalpy Change ($\Delta H$): Defined as the difference between the enthalpies of products and reactants:
$$\Delta H = H_{\text{products}} - H_{\text{reactants}}$$
- Reversibility: If a reaction is reversed, the magnitude of $\Delta H^\circ$ remains the same, but the sign changes.
6.1.6 Conditions for Standard States
- Standard state for a gas is $1 \text{ atm}$.
- Standard state for an element or compound is the most stable physical state at $1 \text{ atm}$ and $25^\circ\text{C}$ ($298 \text{ K}$).
- Standard state for a substance in aqueous solution is $1\text{M}$ concentration.
Relevant Questions & Answers
Q1: Why is energy either evolved or absorbed during a chemical reaction?A: Chemical reactions involve breaking old bonds (consuming energy) and forming new bonds (releasing energy). The net difference between these two processes results in heat being absorbed or evolved.
Q2: Define the term "Thermochemistry."A: Thermochemistry is the branch of chemistry that deals with the heat or thermal energy changes in chemical reactions.
Q3: How does the enthalpy change ($\Delta H$) differ between exothermic and endothermic processes?A: For an exothermic process, $\Delta H$ is negative because heat is released by the system. For an endothermic process, $\Delta H$ is positive because heat is absorbed by the system.
Q4: Interpret the following equation: $H_{2(g)} + I_{2(g)} \rightarrow 2HI_{(g)} \quad \Delta H^\circ = +53.8 \text{ kJ}$A: This shows that when $1 \text{ mole}$ of $H_2$ gas reacts with $1 \text{ mole}$ of $I_2$ gas to give $2 \text{ moles}$ of $HI$ gas at $25^\circ\text{C}$ and $1 \text{ atm}$, $53.8 \text{ kJ}$ of heat is absorbed.
Exercise 6.1: Classification
- (a) Freezing of water: Exothermic.
- (b) Combustion of methane: Exothermic.
- (c) Sublimation of dry ice: Endothermic.
- (d) $H_2O_{(g)} \rightarrow H_2O_{(l)}$: Exothermic.
- (e) Decomposition of limestone: Endothermic.
6.1 Energy in Chemical Reactions
Chemical reactions involve the breaking of old bonds and the formation of new bonds. This process leads to the evolution (release) or absorption of heat energy.
- Bond Breaking: This process always consumes energy.
- Bond Making: This process always releases energy.
- Net Energy Transfer: If energy released by bond formation exceeds energy consumed by bond breaking, heat is transferred to the surroundings. If bond breaking requires more energy than bond formation releases, heat is transferred from the surroundings to the system.
6.1.1 Units of Thermal Energy
Thermal energy is defined as the energy in a system due to the movement of its atoms and molecules.
- Joule (J): The standard SI unit for heat or thermal energy.
- Calorie (cal): The amount of heat required to raise the temperature of $1 \text{ gram}$ of water from $14.5^\circ\text{C}$ to $15.5^\circ\text{C}$.
- Conversion: $1 \text{ Calorie} = 4.18 \text{ Joules}$.
6.1.2 Thermochemical Reactions
Thermochemistry is the branch of chemistry dealing with heat changes in chemical reactions. A Thermochemical Equation is a balanced equation that includes the heat change ($\Delta H^\circ$).
Shutterstock1. Exothermic Reactions
These reactions proceed with the evolution (release) of heat. Energy is transferred from the system to the surroundings, and $\Delta H$ is represented as a negative value.
- Example: $C_{(s)} + O_{2(g)} \rightarrow CO_{2(g)} \quad \Delta H^\circ = -393.5 \text{ kJ}$.
2. Endothermic Reactions
These reactions proceed with the absorption of heat. Energy is transferred from the surroundings to the system, and $\Delta H$ is represented as a positive value.
- Example: $N_{2(g)} + O_{2(g)} \rightarrow 2NO_{(g)} \quad \Delta H^\circ = +180.5 \text{ kJ}$.
6.1.4 Enthalpy Change and Standard States
At constant pressure, heat change is equated to the change in enthalpy ($\Delta H$). It is the difference between the enthalpies of the products and the reactants:
$$\Delta H = H_{\text{products}} - H_{\text{reactants}}$$
6.1.6 Conditions for Standard Heat of Reaction
Standardized $\Delta H$ values are calculated when substances are in their standard state:
- Gases: Pressure of $1 \text{ atm}$.
- Elements/Compounds: Most stable physical state at $1 \text{ atm}$ and $25^\circ\text{C}$ ($298 \text{ K}$).
- Aqueous Solutions: Concentration of $1\text{M}$.
Relevant Questions & Answers
Q1: Why is energy evolved or absorbed in a reaction?A: It is due to the difference between the energy required to break old bonds and the energy released when new bonds are formed.
Q2: What happens to $\Delta H^\circ$ if a thermochemical reaction is reversed?A: The magnitude of $\Delta H^\circ$ remains the same, but the sign changes (e.g., negative becomes positive).
Q3: Define Standard Enthalpy of Combustion ($\Delta H_c^\circ$).A: It is the enthalpy change when one mole of a substance is completely burnt in excess oxygen under standard conditions.
Concept Assessment Exercise 6.1 - Solutions
- (a) Freezing of water: Exothermic.
- (b) Combustion of methane: Exothermic.
- (c) Sublimation of dry ice: Endothermic.
- (d) $H_2O_{(g)} \rightarrow H_2O_{(l)}$: Exothermic.
- (e) Decomposition of limestone: Endothermic.