AP Chemistry 7.1: Dynamic Equilibrium and Unit 7 Review

Chemical equilibrium describes what happens in a reversible reaction when the forward and reverse reactions occur at the same rate. At equilibrium, reactions have not stopped. Instead, reactants are continuously turning into products while products are continuously turning back into reactants.

• This unit is strongly connected to earlier AP Chemistry units. Kinetics explains why equilibrium is dynamic because both forward and reverse reactions continue occurring. Stoichiometry is needed when constructing ICE tables. Gas laws help with equilibrium problems involving partial pressures. Net ionic equations connect directly to solubility equilibria and precipitation.
• The biggest ideas in Unit 7 are understanding what equilibrium actually means, writing and interpreting equilibrium expressions, calculating equilibrium constants, predicting which direction a reaction will move, determining equilibrium concentrations, understanding how a system responds to disturbances, and working with the solubility-product constant Ksp.

Many chemical reactions are reversible, meaning they can proceed in both directions.

• A reversible reaction is written using a double arrow:
• A + B ⇌ C + D

The forward reaction is:

• A + B → C + D
• while the reverse reaction is:
• C + D → A + B
• When a reversible reaction begins with mostly reactants, the forward reaction is usually faster because there are many reactant particles available to collide. As products begin forming, the reverse reaction becomes possible and its rate increases.
• Eventually, the forward and reverse reaction rates can become equal. At that point, the system has reached dynamic equilibrium. Khan Academy’s current Unit 7 defines dynamic equilibrium using equal forward and reverse rates, which results in constant concentrations or partial pressures over time.

At equilibrium:

• rateforward = ratereverse
• This does not mean:
• [reactants] = [products]
• The concentrations can be very different.
• For example, an equilibrium mixture might contain:
• 0.80 M reactant
• and:
• 0.20 M product
• and still be at equilibrium.
• Equilibrium depends on the rates being equal, not the amounts being equal.

Suppose:

• A ⇌ B

At equilibrium, A is still converting into B, and B is still converting into A.

• However, imagine 100 A particles become B during a certain time interval while 100 B particles become A during the same interval.
• There is no net change in either concentration.
• Therefore:
• [A] remains constant
• and:
• [B] remains constant
• The system looks unchanged macroscopically even though reactions continue at the molecular level.
• That is why equilibrium is called dynamic, not static.

Consider beginning with only reactants.

• Initially:
• forward reaction rate = relatively high
• because reactant concentration is high.
• reverse reaction rate ≈ 0
• because almost no products are present.
• As the reaction proceeds, reactants are consumed, so the forward rate generally decreases.
• At the same time, products accumulate, causing the reverse rate to increase.
• Eventually:
• forward rate = reverse rate
• and equilibrium has been reached.

If you graph reaction rate against time, the forward rate may begin high and decrease while the reverse rate begins low and increases.

At equilibrium, the two lines meet.

• After that point:
• rateforward = ratereverse
• Both rates remain nonzero.
• This is one of the best ways to visualize dynamic equilibrium.

On a concentration graph, reactant concentration may decrease while product concentration increases.

• Eventually both curves become horizontal.
• Horizontal curves mean the concentrations are now constant.
• However, the curves do not need to meet.
• If one concentration is higher than another at equilibrium, that is completely normal.

For equilibrium to remain established, substances usually cannot continuously escape from the system.

• For example:
• H₂O(l) ⇌ H₂O(g)
• can establish equilibrium in a sealed container.
• Liquid molecules evaporate while gas molecules condense.
• Eventually:
• rate of evaporation = rate of condensation
• But in an open container, water vapor can escape into the room. That makes it much harder to maintain the same dynamic equilibrium inside the container.

For a reversible reaction:

• reactants ⇌ products
• three situations are possible.
• If:
• rateforward > ratereverse
• there is a net conversion of reactants into products.
• The reaction proceeds to the right.
• If:
• ratereverse > rateforward
• there is a net conversion of products into reactants.
• The reaction proceeds to the left.
• If:
• rateforward = ratereverse
• the system is at equilibrium.

This is probably the biggest misconception in Unit 7.

At equilibrium, particles are still reacting.

• The system is simply converting reactants to products at the same rate that products are converting back to reactants.

Do not say that equilibrium means equal concentrations.

Do not say that reactions stop at equilibrium.

Do not assume equilibrium means there are 50% reactants and 50% products.

Do not confuse constant concentration with zero reaction rate.

At equilibrium:

• forward rate = reverse rate
• and therefore:
• concentrations remain constant
• but the reaction is still happening in both directions.

Kinetics and equilibrium are closely connected but answer different questions.

• Kinetics asks:
• How fast does the reaction happen?
• Equilibrium asks:
• What relative amounts of reactants and products remain once the forward and reverse rates become equal?

At equilibrium:

• rateforward = ratereverse
• This is a kinetic statement.
• But the value of K describes the equilibrium composition.

A catalyst makes both forward and reverse reactions faster, so equilibrium is reached sooner.

• However, the catalyst does not change:
• K
• or:
• equilibrium composition

ICE tables depend heavily on balanced chemical equations.

For:

• N₂ + 3H₂ ⇌ 2NH₃
• if N₂ decreases by x:
• H₂ must decrease by:
• 3x
• and NH₃ must increase by:
• 2x
• These relationships are simply stoichiometric mole ratios applied to concentration changes.

Gas equilibrium often uses partial pressures.

• Remember:
• Ptotal = ΣPi
• and:
• Pi = XiPtotal
• These relationships may be needed before calculating Kp.
• Changes in gas volume also affect equilibrium because they change partial pressures and therefore Q.

Equilibrium: Equal rates do not mean equal concentrations.

Dynamic equilibrium: Reactions continue in both directions.

K expressions: Products go over reactants and coefficients become exponents.

Solids/liquids: Pure solids and pure liquids are omitted from K.

Q vs. K: Q<K means right; Q>K means left.

K magnitude: Large K means product-favored, not fast.

Reaction reversal: Take 1/K.

Multiplying equations: Raise K to the same power.

Adding equations: Multiply the K values.

Temperature: K depends on temperature.

Concentration changes: Changing concentration shifts equilibrium but does not change K at constant temperature.

ICE tables: Changes must follow stoichiometric coefficients.

Small-x: Do not use the approximation unless it is reasonable.

Pressure: Only gaseous species matter when comparing gas moles.

Equal gas moles: Compression may produce no equilibrium shift when both sides contain the same total amount of gas.

Inert gas: Adding one at constant volume generally does not shift equilibrium because reacting-gas partial pressures remain unchanged.

Catalyst: Speeds equilibrium establishment but does not change K or equilibrium composition.

Solids: Adding more pure solid usually does not affect equilibrium if some solid is already present.

Ksp: Never include the solid in the expression.

Molar solubility: Ksp is not automatically equal to s.

Precipitation: Qsp>Ksp means precipitate forms.

Mixing solutions: Recalculate ion concentrations using the total mixed volume before finding Qsp.

Common ion: Usually decreases solubility.

pH: Acid can increase solubility of salts containing basic anions.

Before taking a Unit 7 test, you should be able to explain dynamic equilibrium at both the particle level and macroscopic level. You should understand that forward and reverse reactions continue at equilibrium, that their rates become equal, and that constant concentration does not mean the reaction has stopped.

You should be able to write Kc, Kp, and Q expressions from balanced equations, correctly apply coefficients as exponents, and know which species should be omitted. You should understand the difference between Q and K and confidently predict reaction direction using Q<K, Q>K, or Q=K.

You should be able to calculate an equilibrium constant from equilibrium concentrations or partial pressures and interpret whether a particular K value represents a product-favored or reactant-favored equilibrium.

You should understand the mathematical properties of equilibrium constants: reversing a reaction takes the reciprocal of K, multiplying a reaction by n raises K to the nth power, and adding reactions multiplies their equilibrium constants.

You should be able to construct and solve ICE tables, use reaction stoichiometry in the change row, determine the proper direction of change, solve equations for x, recognize when a quadratic may be required, and judge whether a small-x approximation is reasonable.

You should understand Le Châtelier's principle for changes in concentration, gas volume, pressure, and temperature. More importantly, you should be able to explain those shifts using changes in Q rather than only memorizing rules.

You should know that concentration, pressure, and volume disturbances can change Q without changing K at constant temperature, while temperature changes can change K.

You should understand why catalysts do not shift equilibrium and why adding an inert gas at constant volume usually does not affect equilibrium.

Finally, you should be able to write Ksp expressions, calculate molar solubility, use stoichiometric ion concentrations correctly, predict precipitation using Qsp vs. Ksp, account for dilution when solutions are mixed, explain the common-ion effect, and understand why pH can affect the solubility of certain ionic compounds.

If you can connect dynamic equilibrium → K → Q → ICE tables → Le Châtelier → Ksp, rather than memorizing them as separate ideas, you understand the main logic of AP Chemistry Unit 7.