Kinetics , Thermodynamics and Equilibrium

Kinetics and Thermodynamics Kinetics: deals with rates of reactions (how quickly a reaction occurs) Thermodynamics: involves changes in energy that occur in reactions

Kinetics: Collision Theory Collision theory states that a reaction is most likely to occur if reactant particles collide with the proper energy and orientation. An ineffective collision reaction does not occur An effective collision reaction occurs

Which event must always occur for a chemical reaction to take place? (1) formation of a precipitate (2) formation of a gas (3) effective collisions between reacting particles (4) addition of a catalyst to the reaction system

Factors Affecting Rate

1. Type of substance Ionic substances react faster: bonds require less energy to break AgNO3 (aq)+NaCl(aq)AgCl(s)+NaNO3 (aq) In solution ionic solids dissociate into ions: Ag+ NO3- Na+ Cl- Covalent react more slowly: bonds require more energy to break H2 (g)+I2 (g)2 HI (g) Bonds must be broken then be reformed. (takes more time)

2. Temperature increase Average kinetic energy increases and the number of collisions increases. Reactants have more energy when colliding. This increases rate.

3. Concentration increase Increases rate due to the fact that more particles are in a given volume, which creates more collisions.

4. Surface Area Increase Increases rate due to increased reactant interaction or collisions (powder vs. lump)

5. Pressure Increases Increases the rate of reactions involving gases only As pressure  Volume  so: spaces between molecules  which  the frequency of effective collisions

6. Catalyst Catalyst: substance that increases rate of reaction, provides a shorter or pathway for the reaction to occur. Catalysts remain unchanged during the reaction and can be reused.

Quick Review – Factors that affect reactions Ionic solutions have faster reactions than molecule compounds. (bonding) Temp.  Rate  conc. rate  surface area  rate  Pressure  rate,  P  rate (gas) Catalysts speed up reactions.

What are the two things that must happen in order to have an effective collision?

Potential Energy Diagrams Graphs heat during the course of a reaction.

Exothermic: PE of products is less because energy was lost. PE of reactants (ER) Activation Energy (Ea) PE of Activated Complex PE of products (EP) Heat of reaction (ΔH) = Ep - ER Activation Energy (Ea)* reverse reaction

Endothermic: PE of products is more because energy was gained. PE of products (EP) PE of reactants (ER) Activation Energy (Ea) Heat of reaction (ΔH) PE of Activated Complex Activation Energy (Ea)* reverse reaction

Catalysts

Thermodynamics Heat content (Enthalpy): amount of heat absorbed or released in a chemical reaction Enthalpy (ΔH = Hproducts – Hreactants)

ΔH = PEproducts – PEreactants ΔH is positive when the reaction is endothermic. Heat of products are greater than reactants ΔH is negative when the reaction is exothermic. Heat of reactants were greater than the products

Exothermic Releases heat Endothermic Absorbs energy PE decreases H is negative Energy is on the right of equation 2H2 + O2 2H2O + energy Endothermic Absorbs energy PE increases H is positive Energy is on the left of the equation 2H2O + energy 2H2 + O2 J Deutsch 2003

Table I Includes heats of reaction for combustion, synthesis (formation) and solution reactions. You must remember equation stoichiometry (balanced equations). Endothermic: heat is a reactant (left) Exothermic: heat is a product (right)

Table I- Check for understanding Which reaction gives off the most energy? Which reaction gives off the least energy? Which reaction requires the most energy absorbed to occur?

3

1

2

This is called the Heat of Summation,ΔH AP topic Hess’s Law Hess’s Law states that the heat of a whole reaction is equivalent to the sum of it’s steps. This is called the Heat of Summation,ΔH

Plan a Strategy Evaluate the given equations. Rearrange and manipulate the equations so that they will produce the overall equation. (change coefficients, reverse the reaction). Cross out common reactants/ products in all steps. Change enthalpy (ΔH) terms accordingly and add the (ΔH ) terms. ΔH1 = ΔH2 + ΔH3

For example, suppose you are given the following data: Use the data above to obtain the enthalpy change for the following reaction:

Work Area

This reaction can also be carried out in two steps: Example #1 Our reaction of interest is: N2(g) + 2O2(g) 2NO2(g) ΔH1 = 68 kJ This reaction can also be carried out in two steps: N2 (g) + O2 (g) 2NO(g) ΔH2 = 180 kJ 2NO (g) + O2 (g) 2NO2(g) ΔH3 = -112 kJ

2NO (g) + O2 (g) 2NO2(g) ΔH3 = -112 kJ If we take the previous two reactions and add them, we get the original reaction of interest: N2 (g) + O2 (g) 2NO(g) ΔH2 = 180 kJ 2NO (g) + O2 (g) 2NO2(g) ΔH3 = -112 kJ N2 (g) + 2O2 (g) 2NO2(g) ΔH1 = 68 kJ

Example #2 Using Hess’s Law, determine the ΔH for the following reaction: 4HCl(g) + O2(g) → 2 Cl2(g) + 2H2O(g) This reaction can take place in a series of two steps: H2(g) + Cl2(g) → 2HCl(g) ΔH =- 185 kJ 2H2(g) + O2(g) → 2H2O ΔH = - 483.7 kJ

Work Area

3. Using Hess’s Law, determine the ΔH for the following reaction: 3C(s) + 4H2 (g) → C3H8 (g) This reaction can take place in a series of three steps: C(s) + O2 (g) → CO2 (g) ΔH = -394 kJ C3H8 (g) + 5O2 (g) → 3CO2 (g) + 4H2O(l) ΔH = -2220 kJ H2 (g) + ½ O2 (g) → H2O (l) ΔH = -286 kJ

Entropy (ΔS) Definition: randomness, disorder in a sample of matter Gases have high entropy Solids have low entropy

How to increase entropy: Phase change from s  l  g Increasing the moles of gases Dissolving a substance (going from s to aq) Decomposition reaction

2

2

2

Spontaneous Reactions This is a reaction that will occur naturally. It does not need outside help. Nature favors low energy (more stable) and high entropy

Analogy: Your Bedroom You like to have low enthalpy (low energy) when it comes to household chores. As a result, your room tends to have high entropy (very messy, disorderly). This is what nature prefers: low enthalpy and high entropy.

When DH is - and entropy is + (exothermic) (greater disorder) the reaction would be spontaneous

Are all spontaneous reactions exothermic and with a greater system disorder?

Gibbs free energy is a measure of chemical energy AP topic Gibbs Free Energy Gibbs free energy is a measure of chemical energy Gibbs free energy, ΔG, enables us to predict whether a reaction will be spontaneous

The free energy change is defined as: ΔG = ΔH – TΔS A spontaneous reaction depends on enthalpy, entropy, and temperature

∆G = ∆H - T∆S If a reaction is exothermic (negative ∆ H) and entropy increases (positive ∆S) then ∆G must be NEGATIVE So the reaction is spontaneous

∆G = ∆H - T∆S So the reaction is not spontaneous If a reaction is endothermic (positive ∆H) and entropy decreases (negative ∆S) then ∆Go must be POSITIVE So the reaction is not spontaneous

Gibbs Free Energy, G ∆G = ∆H - T∆S ∆Ho ∆So ∆Go Reaction exo(–) increase(+) – Prod-favored endo(+) decrease(-) + React-favored exo(–) decrease(-) ? T dependent endo(+) increase(+) ? T dependent

Chemical Equilibrium

Reversible Reactions Some chemical reactions are able to proceed in both directions under the appropriate conditions. Example: Fe3O4 (s) + 4 H2 (g) ↔ 3 Fe(s) + 4 H2O(g)

Equilibrium Equilibrium is a dynamic condition where rates of opposing processes are equal. Types of Equilibrium: Phase equilibrium Solution Equilibrium Chemical Equilibrium

Reactions that go to completion do not reach equilibrium: Equilibrium is not reached in an open system . The system must be closed. Gases can escape! Equilibrium is not reached if there is an insoluble product (precipitate)

Phase Equilibrium Rate of phase change is equal when two phases exist at the same temperature. Example: H2O (l)  H2O (g) This happens on the plateau of the heating/cooling curve.

Solution Equilibrium Rate of dissolving = rate of crystallization Occurs in saturated solutions

Chemical Equilibrium Rateforward reaction = Ratereverse reaction Concentration of reactants and products are constant NOT equal.

At equilibrium, the forward and reverse reaction rates are equal.

At equilibrium, the concentrations of the reactants and products are constant

The Concept of Equilibrium RECC That’s what equilibrium means to me!! Rates are Equal Concentrations are Constant

Le Chatelier’s Principle Whenever stress is applied to a reaction at equilibrium, the reaction will shift its point of equilibrium to offset the stress.

Temperature, pressure, changes in reactant or product concentrations Stresses include: Temperature, pressure, changes in reactant or product concentrations

N2 (g) + 3 H2 (g)  2 NH3 (g) + heat The letter T for temperature goes on the side that has heat or kJ. The Letter P goes on the side that has a greater number of gas moles.

The rule for determining which way the reaction shifts is: Add To shifts Away (increase) Take Out shifts Towards (decrease)

N2 (g) + 3 H2 (g)  2 NH3 (g) + heat Concentration increase shift away from increase Concentration decrease shift toward decrease  pressure shifts in direction of fewer gas molecules.  pressure shifts in direction of more gas molecules  temperature favors endothermic reaction Shift away from heat  temperature favors exothermic reaction Shift towards heat

Example: The Haber Process N2 (g) + 3 H2 (g)  2 NH3 (g) + heat  [N2] shift towards products (right)  [H2] shift towards reactants (left)  [NH3] shift towards reactants (left)  [NH3] shift towards products (right)  pressure shift towards products (right)  pressure shift towards reactants (left)  temperature shift towards reactants (left)  temperature shift towards products (right)

Effect of Catalyst: Addition of catalysts changes the rate of both the forward and reverse reactions. There is no change in concentrations but equilibrium is reached more rapidly.

c State the effect on the number of moles of NH3 (g) if a catalyst is introduced into the reaction system. Explain why this occurs.

The Haber Process Application of LeChatelier’s Principle N2 (g) + 3 H2 (g)  2 NH3 (g) + 92 kJ increase pressure Shift  decrease Temp remove NH3 add N2 and H2 ****Maximum yields of NH3 occurs under high pressures, low temperatures and by constantly removing NH3 and adding N2 & H2