Chapter 18: Reaction Rates and Equilibrium

STATE STANDARDS ADDRESSED:

  1. INVESTIGATION AND EXPERIMENTATION: Scientific progress is made by asking meaningful questions and conducting careful investigations. As a basis for understanding this concept and addressing the content in the other four strands, students should develop their own questions and perform investigations.
  2. REACTION RATES: Chemical reaction rates depend on factors that influence the frequency of collision of reactant molecules.
  3. CHEMICAL EQUILIBRIUM: Chemical equilibrium is a dynamic process at the molecular level.
  4. CHEMICAL THERMODYNAMICS: Energy is exchanged or transformed in all chemical reactions and physical changes of matter.


18.1 Reaction Rates and Collision Theory
1. Rates measure changes that occur within intervals of time.
2. Chemical rates are expressed using chemical units and time. Eg., moles/s.
3. Collision theory says that atoms, ions, and molecules can form a chemical bond when they collide, provided the particles have enough kinetic energy.
4. Particles without enough energy still collide but do not react.

Figure 18.2

5. The bonds holding molecules together can break apart.
6. The minimum energy colliding particles must have in order to react is the activation energy.
7. An activated complex is the arrangement of atoms at the peak of the activation energy barrier.
8. The activated complex is sometimes called the transition state. It is unstable and as likely to form reactants or products.
9. The lifetime of a complex is about 10-13 seconds.
10. It is a combination of reactants before reaching final form (transition).

Figure 18.3

11.Some spontaneous reactions are slow at room temperature. E.g., C and O combine spontaneously; however,  at room temp. activation energy barrier is not overcome enough to break C-C and O-O bonds. Reaction rate is zero.
18.2 Factors affecting Reaction Rates
1. Each chemical reaction proceeds at its own rate.
2. Rate of most reactions can be modified in several ways, explained by collision theory.
Effect of Temperature. Speeds up reactions. Increase in temperature increases number of particles with KE to react when they collide. Frequency of collisions increases. More molecules cross activation energy barrier. Products form faster.

E.g., charcoal does not burn at room temp. After starter flame, some C and O react to form CO2. This reaction releases energy to cause others to form, etc. Reaction continues after activation energy removed, until all C and O gone.

Effect of Concentration. Number of reacting particles in given volume affects rate. More increases collision frequency. E.g., burning splint placed in pure O2 flames due to increased concentration.
Effect of Particle Size. The smaller the particle size, the greater the surface area for a given mass. Increases collision frequency and rate of reaction.

Effect of Catalyst. A catalyst is a substance that increases the rate of a reaction without being used up itself in the reaction. A catalyst lowers the activation energy barrier. More reactants can form products in a given time.

Figure 18.7

Few reactions in body (37oC) would occur fast enough without catalysts called enzymes. (Protein in stomach would take 50 years to digest)
3. An inhibitor is a substance that interferes with the action of a catalyst. (Aspirin).
18.3 Reversible Reactions
1. Not all reactions go completely to products.
2. Some reactions are reversible.

3.In reversible reactions the reactions in both directions occur simultaneously - Forward reaction and Reverse reaction.
4. Oxygen and sulfur dioxide is used up and sulfur trioxide is formed.
5. As sulfur trioxide is built up, oxygen and sulfur dioxide begin to form in reverse. As sulfur trioxide becomes higher, the reverse reaction speeds up.
6. Eventually chemical equilibrium is reached when forward and reverse reactions are taking place at the same rate.
7. This does not necessarily mean concentrations of both are same.
8. The equilibrium position of a reaction is given by the relative concentrations of the components of the system at equilibrium.

9. If reaction uses up reactants - goes to completion or is irreversible.
10. If reaction forms no products - say no reaction.

11. Catalysts speed up forward and backward reactions equally. Decrease time for equilibrium to be established.
18.4 Factors Affecting Equilibrium
1.Equilibrium is a delicate balance. Changes disrupt balance. Systems make minimum adjustment.
2. There may be a shift in amount of reactants or products called shift in position of equilibrium.
3. Equilibrium shift compensates for the disturbance that caused it.
4. Changes in concentration:

5. Adding a product always pushes a reversible reaction in the direction of reactants.
6. Removing a product always pulls a reversible reaction in the direction of products.
7.Removing product is a trick to increase yield of reaction.
8. When a reactant is added to a system at equilibrium, the reaction shifts in the direction of the formation of products. When a reactant is removed, the reaction shifts in the direction of formation of reactants.
9. Changes in Temperature:
10. Increasing the temperature causes the equilibrium position of a reaction to shift in the direction that absorbs heat.
11. Heat can be considered a product or reactant.


(Exothermic Reaction)

12. Adding or removing heat from either side can force reaction towards a new equilibrium.
13. Changes in Pressure - Affects only equilibrium of a system with unequal number of moles of gaseous reactants or products.

14. If pressure in a cylinder is increased, increasing density, then the molecules above will form (ammonia) to reduce number of molecules is system. Pressure will not decrease to original pressure, but will decrease. Equilibrium favors more product.
15. If reduce pressure (pull piston up) reaction favors reactants.

Figure 18.13

16. Le Chatelier's principle: If a stress is applied to a system in a dynamic equilibrium, the system changes to relieve the stress.
17. Stresses that upset equilibrium of a chemical system: Changes in concentration, changes in temperature, changes in pressure.
18. Do Example 1.
18.5 Free Energy
1. Free energy is energy that is available to do work. Energy is released from chemical and physical processes.
2. Free energy may be available but is not always used efficiently.
3. Auto engine uses only 30% of free energy of burning gasoline. 70% is lost as friction and heat.
4. Living organisms are better; however, not more than 70% efficient. This can be useful in itself as body temperature (37oC) is maintained by wasted heat.
18.6 Spontaneous Reactions
1. Though you can write a balanced equation, it may not represent what will really happen.

E.g.: CO2(g) ----- C(s) + O2 (g)

2. This is reverse of combustion reaction and is a decomposition reaction.---- does not occur.
3. Spontaneous reactions naturally favor the formation of products at specified conditions.
4. All spontaneous reactions release free energy and are said to be exergonic.
4. Nonspontaneous reactions do not favor the formation of products at the specified conditions.Nonspontaneous reactions do not give substantial products at equilibrium (more than 50% reactants to products).

5. Forward reaction, above, is spontaneous and releases free energy. Reverse is not.

(See 2nd reaction, page 505.)

6. In reversible reactions, one of the reactions is always spontaneous and the other is always nonspontaneous.
7. Spontaneous reactions do not refer to speed. Some spontaneous reactions are slow.
8. E.g., combustion of sugar and oxygen in air will take thousands of years. With heat it goes faster, but still it is spontaneous.
9. Reactions which are nonspontaneous may be made spontaneous under different conditions: temperature, pressure.
10. Nonspontaneous reactions can be coupled to spontaneous reactions. Occurs often in living systems: such as, the energy released from glucose to drive growth, etc.
18.7 Entropy
1. Enthalpy (ch 10) is the amount of heat that a substance has at a given temperature and pressure.
2. Many reactions are spontaneous and give off heat (exergonic). Sometimes not so. E.g., water melts spontaneously at 25 oC and absorbs 6.0 kJ/mole.
3. Entropy is a measure of the disorder of a system. (Scattered marbles have a higher entropy than gathered).
4. The law of disorder states that things move in the direction of maximum disorder or randomness (chaos).

For a given substance entropy of gasliquidsolid
Entropy increases when solids go to liq or gas or liq goes to gas.
Entropy increase when substances divide into parts: NaCl dissociates in water.
Entropy increases in chemical reactions when number of product moleculesreactant molecules.

18.8 Heat, Entropy, and Free Energy
1. The size and direction of heat (enthalpy) changes and entropy changes determine whether reaction is spontaneous.
2. If both exothermic and increase in entropy, then spontaneous.
3. If reduction of entropy is offset by heat given off, then spontaneous.
4. If endothermic is offset by increase in entropy, then spontaneous.
5. Ice melting is endothermic. Favorable entropy change.
18.9 Equilibrium Constants
1. Equilibrium constants relate the amounts of reactants to products at equilibrium.
2. The equilibrium constant Keq is the ratio of product concentrations to reactant concentrations, with each concentration raised to a power given by the number of moles of that substance in the balanced chemical equation.

3. [ brackets] = concentration M/l
4. Value of constant will change with temperature.
5. If Keg 1, products are favored.
6. If Keq < 1, reactants are favored = non spontaneous.
7. Example 2, 3, 4; problems 20 - 24.
18.10 Entropy Calculations
1. Entropy is a quantitative measure of the disorder of a system. Symbol = S. Units are J/K or J/(K x mol) for a mole  or J/K-mol.
2. Theoretical entropy of a perfect crystal at 0 K is zero. Change in entropy is products minus reactants
3. See table 18.2 for entropies.DSo
4. Example 5, 6; problem 26.
18.11 Free Energy Calculations
1. The Gibbs free energy change,DG is the maximum amount of energy that can be coupled to another process to do useful work.
2. Use Table 18.4 to determine  DGo (standard free energy change) = DGof(products)- DGof (reactants)
3. negative answers indicate exergonic reaction and spontaneous.
4. Example 7, 8,
5. For example 8, recall that

DGo = DHo - DTSo.

T is temperature Kelvin.
6. Remember to place units in kJ/mol to make same as enthalpy change.
18.12 Rate Laws
1. The rate of a reaction depends in part on the concentrations of the reactants.
2. Rate = k x [A].
3. A rate law is an expression that relates the rate of a reaction to the concentration of REACTANTS. The specific rate constant for a reaction (k) is a proportionality constant relating the concentrations of reactants to the rate of the reaction.
4. In a first-order reaction the reaction rate is directly proportional to the concentration of only one reactant.
18.13 Reaction Mechanisms
1. A reaction progress curve includes all changes in a reaction.
2. An elementary reaction is one when reactants are converted to products in a single step. One activated complex.
3. Reaction mechanism includes all of the elementary reactions of a complex reaction.
4. progress curve resembles a number of hills and valleys.
5. An intermediate is a product of a reaction that becomes a reactant of another reaction within a reaction mechanism.

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