CHM110H5 Chap.6 Equilibrium and Le Chatelier's Principle
Equilibrium and Le Chatelier's Principle
What equilibrium means in this course
Unit 2 of CHM110H5 is built around chemical equilibrium. You are expected to define the equilibrium state, explain how temperature, pressure, concentration and solvent affect the extent of a reaction, and apply those ideas to reactions in both the gas phase and solution.
Equilibrium is dynamic: the forward and reverse reactions keep running, but at equal rates, so the amounts of each species stop changing. The equilibrium constant K captures where that balance sits. A large K tells you products dominate once the system settles, and a small K tells you very little reactant converts.
Because K depends only on temperature, any mixture of the same reaction at the same temperature ends up with the same ratio, wherever it started.
Calculating constants and concentrations
The learning outcomes ask you to calculate equilibrium constants and equilibrium concentrations by several methods, and to decide when simplified calculations are valid.
The reaction quotient Q has the same form as K but uses the concentrations you have right now. Comparing them gives the direction of change: if Q is smaller than K, the reaction proceeds forward; if Q is larger, it runs in reverse. For numerical problems, an ICE table (initial, change, equilibrium) turns the situation into one unknown.
When K is small relative to the starting concentration, you may drop x where it is subtracted from a large value, but you must then confirm that x is under 5 percent of that value. If the check fails, solve the quadratic instead. Pure solids and liquids such as water do not appear in K.
Le Chatelier's principle in the lab
Experiment 3 puts these ideas in your hands with sparkling water.
Carbon dioxide gas dissolves to form aqueous carbon dioxide, which reacts with water to release hydrogen ions and bicarbonate, so the solution is weakly acidic. Bromocresol green indicator shows the pH inside a capped syringe. Pulling the plunger out expands the gas space and lowers its pressure, so carbon dioxide escapes from solution and the pH rises. Pushing it in raises the pressure, more gas dissolves, and the pH falls.
Venting the syringe removes gaseous carbon dioxide; drawing in fresh gas from a cylinder adds it back.
A warm water bath and an ice bath test the temperature effect: because dissolving carbon dioxide releases heat, warming lowers its solubility and cooling raises it.
The titration and the report
Part B measures dissolved carbon dioxide directly by titrating a 10 ml sample of sparkling water with standardized sodium hydroxide of about 0.1 M, using phenolphthalein to spot the faint pink end point.
Carbon dioxide and hydroxide react one to one to form bicarbonate, so moles of base used equal moles of carbon dioxide in the sample. Marking feedback on this report repeatedly points out two things: dilution formulas do not apply to a neutralization, and every conclusion should state clearly which way the equilibrium moved and which species, gaseous or aqueous carbon dioxide, you are describing.
What this chapter covers
- 01
Dynamic equilibrium and the meaning of K
- 02
Reaction quotient Q versus K
- 03
Kp, Kc and pure solids or liquids
- 04
ICE tables and the 5 percent test
- 05
Le Chatelier shifts: concentration, pressure, temperature, catalyst
- 06
Henry's law and dissolved carbon dioxide
- 07
Experiment 3 syringe and NaOH titration
Worked example · free
Deciding direction, then solving with an ICE table
- 1Initial Q = [H2][I2]/[HI]^2 = (0)(0)/(0.400)^2 = 0, which is less than Kc = 0.020, so the net reaction runs forward, producing H2 and I2.
- 2ICE table: I: [HI] = 0.400, [H2] = 0, [I2] = 0. C: -2x, +x, +x. E: 0.400 - 2x, x, x. So Kc = x^2/(0.400 - 2x)^2 = 0.020.
- 2Take the square root of both sides: x/(0.400 - 2x) = sqrt(0.020) = 0.1414. Then x = 0.05657 - 0.2828x, so 1.2828x = 0.05657 and x = 0.0441 M. No approximation is needed.
- 2Equilibrium values: [H2] = [I2] = 0.0441 M; [HI] = 0.400 - 2(0.0441) = 0.312 M. Check: (0.0441)^2/(0.312)^2 = 0.00194/0.0973 = 0.0200, matching Kc.
- 1Compression doubles every concentration. Both sides contain 2 mol of gas, so Q = (2)(2)/(2)^2 = 1 times its old value, still equal to Kc: no shift occurs.
Key terms
- Equilibrium Constant
- The ratio of product to reactant concentrations, each raised to its coefficient, once forward and reverse rates are equal; it changes only when temperature changes.
- Reaction Quotient
- The same ratio as K but evaluated with the current concentrations; comparing Q with K tells you whether a mixture reacts forward, in reverse, or not at all.
- ICE Table
- A three-row layout of initial concentrations, changes written in terms of x using the coefficients, and equilibrium expressions that are substituted into K.
- Five Percent Test
- A validity check for dropping x: the approximation stands only if x is below 5 percent of the initial concentration it was subtracted from.
- Le Chatelier's Principle
- A system at equilibrium responds to an imposed change in concentration, pressure or temperature by shifting in the direction that partly counteracts that change.
- Henry's Law
- The concentration of a dissolved gas is proportional to its partial pressure above the liquid, with a proportionality constant that falls as temperature rises for carbon dioxide.
Equilibrium and Le Chatelier's Principle FAQ
Does adding a catalyst change the equilibrium constant?
No. A catalyst lowers the activation barrier for the forward and reverse reactions by the same factor, so equilibrium is reached sooner but the final composition and the value of K are identical to the uncatalysed system.
Why does warm sparkling water go flat faster than cold?
Dissolving carbon dioxide gas in water releases heat. Raising the temperature therefore favours the reverse process, the Henry constant shrinks, and less gas stays in solution, so bubbles leave the liquid more quickly and the pH climbs.
What should I do if the small-x approximation gives an answer over 5 percent?
Discard the shortcut result, rearrange the K expression into a quadratic of the form ax^2 + bx + c = 0, keep the positive physically sensible root, and substitute all equilibrium values back to confirm they reproduce K.
Exam move
Work through at least five ICE problems in which you compute Q first and write the direction before touching x. For each, run the 5 percent test out loud and note the ratio of initial concentration to K. Build a one-page table of stresses with columns for shift direction and effect on K, then cover the answers and quiz yourself.
Before the Experiment 3 report, rewrite each syringe observation as a sentence naming CO2(g) or CO2(aq), the shift direction, and the pH change.
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