Chapter 17 Equilibrium

17.1 How Chemical Reactions Occur

Particle collisions

1. Particles MUST collide for a chemical reaction to occur.

2. Orientation of particle collision is important.

3. Effectiveness of particle collision is important.

Figure 17.2: (a)Two BrNO molecules approach each other at high speeds. (b) The

collision occurs. (c) The energy of the collision causes Br-N bonds to break and Br- Br bonds to

form. (d) The products: one Br2 and two NO molecules.

Figure 17.3: Reaction progress of two BrNO molecules.

17.2 Conditions that Affect Reaction Rates

Concentration- Increasing concentration will increase the number of particles that are present, allowing for more collisions

Temperature- increasing the temperature gives the particles more kinetic energy; the collisions, when they occur, will be more effective

Catalyst- a catalyst makes a chemical reaction occur faster (orientation of reacting particles) and is not consumed during the reaction. A catalyst is recoverable (write it over the arrow)

Enzymes- an organic catalyst (catalase, lipase, diastase)

Energy diagram (catalyzed & uncatalyzed)(note especially the difference between exothermic and endothermic reactions)

Figure 17.4: Comparison of the activation energies for an uncatalyzed reaction and for

the same reaction with a catalyst present.

17.3 Heterogeneous Reactions

Homogeneous- all substances are in the same phase of matter

Example:

N2(g) + 3H2(g) 2NH3(g)

Heterogeneous- substances are in different phases of matter

Example:

CaCO3(s) CaO(s) + CO2(g)

17.4 The Equilibrium Condition

17.5 Chemical Equilibrium: A Dynamic Condition

Physical equilibrium (see demo in class: vaporization & condensation, discussed in Chapter 14)

Chemical equilibrium- a dynamic state where the concentrations of all reactants and products remain constant

The rate of the forward reaction equals the rate of the reverse reaction. All chemical species will be present at the same time. Technically all chemical reactions are reversible (although, many of them do not do this spontaneously; they require too much energy input to react in the reverse direction.)

Example:

2Mg(s) + O2(g) 2MgO(s)

2Na(s) + Cl2(g) 2NaCl(s)

Figure 17.8: (a) Net transfer of molecules from the liquid state to the vapor state. (b) The amount of the substance in the vapor state becomes constant. (c) The

equilibrium state.

Figure 17.9: (a) Equal numbers of moles of H2O and CO are mixed in a closed container. (b) The reaction begins to occur. (c) The reaction continues, and more

reactants are changed to products. (d) No further changes are seen as time continues to pass.

Figure 17.10: Changes with time in the rates of reactions.

17.6 The Equilibrium Condition: An Introduction

17.7 Heterogeneous Equilibria

For the generic reaction:

aA + bB cC + dD

Equilibrium expression- temperature defined; no unit

K =

Do not include pure solids or pure liquids (only gases & solutions)

Large value of K- products are favored

Small value of K- reactants are favored

Example: Write the equilibrium expressions for the following reactions

Problem:

Problem:

17.8 Le Chatelier’s Principle

When a change is imposed on a system at equilibrium, the position of the equilibrium shifts in a direction that tends to reduce the effect of the change

Temperature

Problem:

Pressure

Change in volume

Problem:

Increase # moles of reactant

Decrease # moles of product

For example, in the Haber process:

N2(g) + 3H2(g) 2NH3(g)

For example, in the Haber process:

N2(g) + 3H2(g) 2NH3(g)

Figure 17.11: (a) Initial equilibrium mixture. (b) Addition of N2.

(c) New equilibrium position.

Figure 17.12: System initially at equilibrium.

Figure 17.12: Piston is pushed

in.

Figure 17.13: (a) A mixture of NH3(g), N2(g), and H2(g) at equilibrium. (b) The volume is suddenly decreased. (c) The

new equilibrium position.

17.9 Applications Involving the Equilibrium Constant

17.10 Solubility Equilibria

Because: AgCl(s) Ag+(aq) + Cl-(aq)

And Ag+(aq) + Cl-(aq) AgCl(s)(precipitates easily)

Therefore, it is an equilibrium system:

AgCl(s) Ag+(aq) + Cl-(aq)

Ksp

(do not use pure solids in an equilibrium expression)

Example:

Working backwards (know value of Ksp, determine solubility in g/ml or mole/liter)(This is the most difficult math in the entire course!)Example:

Common Ion EffectRemember LeChatelier’s Principle: the presence of a common ion may shift a system in equilibriumWater softening systems (sodium replacement)

Diprotic & triprotic acid dissociation (presence of hydrogen ion actually inhibits further dissociation (common ion).

More complex systems (estuaries, kidney dialysis, etc.)

Weak Acid equilibria, Ka

Weak Base equilibria, Kb

Water equilibrium, Kw