Reaction Energy, Kinetics & Equilibrium

Thermochemistry

• Thermochemistry involves the study of energy transfers as heat that accompany chemical reactions.

Calorimetry and Enthalpy

• A calorimeter is used to measure the amount of energy absorbed or released in a reaction.

• The change in heat is represented by ΔH (enthalpy).

• ΔT represents the temperature change from the beginning to the end of a reaction.

• This helps classify reactions as endothermic or exothermic.

Endothermic Reactions

• Endothermic reactions absorb energy (reactants).

• Bonds are broken, and the surroundings get cold.

• General form: A + B + Energy → AB

Exothermic Reactions

• Exothermic reactions release energy (product).

• Bonds are made, and the surroundings get hot.

• General form: ABC → A + B + C + Energy

Temperature and Kinetic Energy

• Temperature is the average kinetic energy of all particles in a reaction.

• Energy is always involved (lost or gained) through phase change transformations.

Measurement of Heat Energy

• Temperature or heat energy is measured in Joules (J) or Kilojoules (KJ).

• The main formula for calculation is: Q = mCΔT

  • Q = Heat energy

  • m = mass in grams

  • C = Specific heat capacity (from Table B, e.g., 4.18 for water)

  • $ΔT$ = Change in temperature

Energy Review Example:

• How much heat energy is needed to raise 20 grams of water from 10°C to 40°C?

  • Q = ?

  • m = 20 grams

  • C = 4.18 (Table B)

  • $ΔT$ = 40 - 10 = 30

  • Q = mCΔT = (20)(4.18)(30) = 2508 J or 2.5 KJ

Enthalpy Sign Conventions (Table I)

• $ΔH$ (-) = Exothermic (Release Heat)

• $ΔH$ (+) = Endothermic (Absorb Heat)

• Table I indicates whether reactions are endothermic or exothermic based on the sign of $ΔH$.

• It's crucial to quantify the energy released or absorbed based on the number of moles of reactants and products.

Calculating Enthalpy Change

• $ΔH$ = Change in Heat = Enthalpy

• Mathematically, $ΔH$ = [Heat of Products] – [Heat of Reactants]

• A negative value indicates an exothermic reaction, while a positive value indicates an endothermic reaction.

• Knowing this is crucial for chemists planning experiments.

Using Table I ΔH Values

• Table I provides $ΔH$ values for reactions under standard conditions.

• Example:

  • $2CO (g) + O<em>2 (g) → 2CO</em>2 (g)$, $ΔH$ = -566.0 kJ (Exothermic)

  • $2NO (g) → N<em>2 (g) + O</em>2 (g)$, Written backwards: $ΔH$ = -182.6 kJ (Exothermic)

Table I Example Questions

• In the reaction $2CO (g) + O<em>2 (g) → 2CO</em>2 (g)$, how much energy is released if 1.5 moles of $O_2$ are used?

  1. Look up the reaction (written the same or backwards).

  2. Note the amount of energy stated.

  3. Set up a proportion and solve:

     • $1 (O<em>2) / -566 kJ = 1.5 (O</em>2) / X kJ$

     • X = -849 kJ

• How much energy is absorbed when 0.2 moles of $H<em>2$ react in the reaction $H</em>2 (g) + I_2 (g) → 2 HI (g)$

  • $1 (H<em>2) / 53 kJ = 0.2 (H</em>2) / X kJ$

  • X = 10.6 kJ

Chemical Kinetics

• Studies the rate at which a chemical reaction occurs and the mechanism or steps it takes to complete.

  • Rate = Speed

  • Mechanism = Direction

• Reactions can be fast, slow, forward, or reverse.

Kinetics and Kinetic Molecular Theory (KMT)

• Kinetics involves moving particles, so understanding particle behavior in different states of matter is important.

• KMT: All particles of matter are in constant random motion.

• When particles move/collide, the speed at which the reaction takes place increases.

Collision Theory

• Effective collisions occur when collisions are frequent and effective with:

  • Enough Energy

  • Right Orientation

• Collisions have a direct impact on how fast a reaction goes, so chemists can affect the rate of reactions.

Factors Affecting Reaction Rates

1. Increase Temperature

2. Agitation

3. Increase Surface Area

4. Increase Pressure on (g)

5. Nature of Reactants

6. Catalyst: Lowers Activation Energy & Provides an Alternative Pathway

Reaction Mechanism

• Step-by-step process of a reaction occurring in the forward or reverse direction.

• Analyzed using Potential Energy (P.E.) Diagrams.

• These diagrams illustrate how reactions take place regarding energy changes.

Potential Energy Diagram Vocabulary

1. P.E. of Reactants: The amount of energy the reactants contain.

2. Activation Energy: The amount of energy needed to start a reaction.

3. Activated Complex: The maximum amount of energy the reaction can contain.

4. P.E. of Products: Potential Energy of the products.

5. Catalyst: Speeds up a reaction in both directions.

6. $ΔH$ or Enthalpy: Change in Heat between the Products and Reactants. $[Products] – [Reactants]$

Endothermic Reaction Diagram

• ABC + Energy → A + B + C

• $ΔH$ = (+)

Exothermic Reaction Diagram

• A + B + C → ABC + Energy

• $ΔH$ = (-)

Le Chatelier’s Principle & Equilibrium

• Le Chatelier’s Principle: Any reaction in a closed system, when stressed, can correct itself, maintaining a balance called Equilibrium.

At Equilibrium:

1. Reaction rates are equal.

2. Concentrations are constant.

3. Catalysts have no net effect.

   • Reactants are created to Products

   • No overall change

Stressors of a Reaction

1. Concentration or Heat:

   • Adding: The reaction shifts opposite.

   • Removing: The reaction shifts to the same side.

2. Pressure:

   • ↑ Pressure: The reaction shifts to the smaller side or ↓ volume (g) only # of moles determines shift

   • ↓ Pressure: The reaction shifts to the bigger side or ↑ volume

Golden Rule

• Whichever side the reaction shifts to, the concentrations increase.

• Example: $N<em>2 (g) + 3H</em>2 (g) → 2NH_3 (g) +91.8kJ$

  • ↑ [NH3] shifts to N2

  • ↓ [N2] shifts to N₂

  • ↑ [H2] shifts to H₂

  • ↓ [NH3] shifts to NH3

  • ↓ [H2] shifts to Energy

Entropy (ΔS) and Free Energy (ΔG)

• $ΔG^º=ΔH^º - TΔS^º$

Enthalpy, Entropy, and Gibbs Free Energy Overview

• $ΔH$ = (-) = Exothermic (release of Heat)

• $ΔH$ = (+) = Endothermic (absorb Heat)

• To understand reactions fully, we need to know Entropy and Gibbs Free Energy.

Entropy Defined

• Entropy (ΔS) is a measure of the degree of randomness of particles or molecules in a system.

• As Kinetic Energy Increases, Entropy Increases

Entropy in Different States of Matter

• Entropy increases from solids to liquids to gases.

Factors Affecting Entropy

1. ↑ Temp. ↑ Particle Spacing

2. ↓ Pressure ↑ Particle Spacing

3. ↑ Volume ↑ Particle Spacing

4. ↑Surface Area ↑ Particle Spacing

Chemical Processes in Nature

• Driven in two directions:

  1. ↓ Enthalpy $ΔH$ = (-)

  2. ↑ Entropy $ΔS$ = (+)

• When these oppose each other, the dominant factor will dominate for a given system.

• Free Energy (𝝙G)

  • Only the Change in Free Energy can be measured.

Free Energy Change (ΔG)

• At constant pressure and temperature, the Free – Energy change ΔG, is defined as: The difference between the change in Enthalpy, ΔH, and the product of the Kelvin temperature and the Entropy change, ΔS Using the expression:

• $ΔG^o = ΔH^o - TΔS^o$

• Using this formula, you can also calculate ΔH, T, or ΔS as long as the other variables are provided.

Example:

• For the reaction: $NH<em>4Cl (s) → NH</em>3 (g) + HCl (g)$, at 298.15 K, $ΔH = 176 kJ/mol$ and $ΔS = 0.285kJ/(mol K)$.

• Calculate ΔG, and tell whether this reaction can proceed in the forward direction at 298.15 K.

Other Examples:

• Based on the following values, compute ΔG or any other value based on the other information provided:

  • $ΔG = ?$

  • Would it be spontaneous?

  • $ΔH = -85.2 kJ/mol$

  • $T = 127^oC$

  • $ΔS = 0.125 kJ/(mol•K)$

• $ΔG = -237.2 kJ$

• $ΔS = 0.465 kJ/(mol•K)$

• What is AH?

• $T = 400 K$

• $ΔG = ?$

• Would it be spontaneous?

• $ΔH = + 38.6 kJ/mol$

• $T = 300^oC$

• $ΔS = 0.243 kJ/(mol•K)$

Equilibrium Expression (Keq)

• Also called “Mass Action Expression”

• Relates the concentration of products to reactants once equilibrium has been reached.

• For this general reaction: $aA + bB ↔ cC + dD$

• $Keq = {[C]^c x [D]^d} / {[A]^a x [B]^b}$

• “[ ]” the brackets mean “the concentration of”

Equilibrium Expression Details

• Concentration of products over concentration of reactants raised to the power of their coefficient in balanced equation

• $Keq = {[C]^c x [D]^d} / {[A]^a x [B]^b}$

• **IMPORTANT ** Excludes: solids and pure liquids as they do not have any effect on concentration values.

Equilibrium Expression Example

• Ex: Write Keq expression for: $N<em>2(g) + 3H</em>2(g) ↔ 2NH_3(g)$

• All gases (nothing excluded)

• $Keq = {[NH<em>3]^2} / {[N</em>2] x [H_2]^3}$

Equilibrium Expression Example with Exclusion

• Ex: Write Keq expression for: $2NO(g) + 2H<em>2(g) ↔ N</em>2(g) + 2H_2O(l)$

• $Keq = {[N<em>2]} / {[NO]^2 x [H</em>2]^2}$

• Take out pure liquid