CHEM 1422 Exam 1

Chapter 15 and 17.4

Equilibrum

_ is when two equal and opposite actions occur at the same time so there is no apparent change in the system

Dynamic Equilibrium

E.g., Water pours into a bucket; water flows out of the bucket through a hole in the bottom. Here, the water level will reach the so-called _ if the operations of fill and empty occur at the same rate

Reversible

Reactions which simultaneously occur in both the forward and the reverse directions, so there is always a non-zero concentration of both the products and the reactants in the reaction mixture, are called _

Closed System

nothing is allowed to escape the system

as fast as

In chemistry, a reversible chemical reaction in a closed system (nothing is allowed to escape the system) can reach a dynamic equilibrium if the reactants are sued up to form the products _ the products decompose back to the reactants

completion

Fact: very few reactions run to _, i.e. where reactants can only get converted into products, and no products can decompose back into the reactants → Most reactions are reversible, and at one point an equilibrium (=m) establishes

Solubility reactions:

CaCO₃(s) ← → Ca²⁺(aq) + CO₃^2-(aq)

Weak acid ionization reactions:

HF(aq) ← → H⁺(aq) + F^-(aq) or HF(aq) + H₂O(l) ← → H₃O⁺(aq) + F^-(aq)

Weak base ionization reactions:

NH₃(aq) + H₂O(l) ← → NH₄⁺(aq) + OH^-(aq)

Complex formation and dissociation reactions:

Fe³⁺(aq) + 6 H₂O(l) ← → Fe(H₂O)₆³⁺(aq)

Gas equlibria:

2 NO₂(g) ← → N₂O₄(g)

EQUILIBRIUM CONSTANT, K

It has been found that, when a reversible chemical reaction comes to equilibrium… no matter what the initial concentrations are, the same special ratio of equilibrium concentrations is always achieved… as long as the temperature stays the same… and we call this ratio the _

K

Each reversible process has its own value of _

temperature

K changes with _

concentrations (K_c), partial pressures (K_p), or both (K, aka thermodynamic equilibrium constant),

K can be expressed in terms of _ and has no units

K_c = [C]_eq^c[D]_eq^d / [A]_eq^a[B]_eq^b

For a general equation: aA + bB ← → cC + dD

reaction quotient, Q, expression

Expression similar to the equilibrium constant expression exists also for concentrations that are not necessarily at =m, and is called a _

forward (toward the products)

If Q < K the reaction proceeds _

is at equilibrium

If Q = K the reaction _

in reverse (toward the reactants)

If Q > K the reaction proceeds _

one and the same ratio

Note: regardless of the initial concentrations, at a given temperature, one and the same ratio of products to reactants is established when the equilibrium is reached

product-favored

if K » 1; the reaction is _ → there is much more product than reactant at =m so we say: “=m lies to the right”

reactant-favored

if K >« 1; the reaction is _ → there is much more reactant that product at =m so we say: “=m lies to the left”

pressures

An equilibrium constant can also be expressed in terms of _: For a general equation: aA + bB ← → cC + dD. At =m, K_p = P(C)_eq^c P(D)_eq^d / P(A)_eq^a P(B)_eq^b

K_c(RT)^Δn_gas

K_p = _

K_p(RT)^-Δn_gas

K_c = _

0.082057 L atm / mol K

It can be shown using the ideal gas law (PV = nRT), that: K_p = K_c(RT)^Δn_gas and K_c = K_p(RT)^-Δn_gas where R = _, P (atm), and T(K)

Δn_gas

_ = n_gas(products) - n_gas(reactants)

are the same

Note: for reactions where Δn_gas = 0, K_c and K_p _

Homogeneous Equilibria

_ = equilibria in which all reactants and products are in the same phase (s, l, g or aq) e.g., N₂O₄(g) ← → 2 NO₂(g)

Heterogeneous Equilibria

_ = equilibria involved reactants and products in different phases

thermodynamic

Here, most commonly (unless instructed otherwise), we construct the _ equilibrium constant, K, in which we* use pressures for gases and molarities for solutions and we omit concentration terms for pure solids and liquids

temperature

Rules for K: (i) All =m constants change with _

independent

Rules for K: (ii) =m constant is _ of initial concentrations

reciprocal

Rules for K: (iii) If the original reaction is reversed, the new =m constant is the _ of the original one

raised

Rules for K: (iv) If the original equation is multiplied by a number, the new =m constant is the original =m constant _ to the same number

product

Rules for K: (v) If two or more reactions are added, the new =m constant (i.e., the =m constant for the sum of the reactions) is the _ of the original =m reactions

RECONSTITUTE

WHILE EQUILIBRIUM CONSTANTS HAVE NO UNITS, WHEN INTERPRETING RESULTS MAKE SURE YOU “_” CORRECT UNITS FOR PRESSURE OR CONCENTRATION

Solubility Product Constant, K_sp

An =m establishes between an undissolved solid and ions in solution. Insoluble ionic compound (s) ← → ions(aq), K_sp «1

poorly soluble

Fact: some ionic salts, some metal hydroxides and some metal oxides are _ (aka insoluble)

Molar Solubility, s

_ = number of moles of the solute that dissolve in forming 1 L of a saturated solution (units: mol/L; often, the solvent is WATER)

Saturated solution

_ is one in which the solution is in contact with undissolved solute. Adding ions to the saturated solution would cause more solid to precipitate.

different stoichiometries

If ionic compounds dissociate according to _, e.g., 1:1, 1:2, 1:3, 2:3, etc…, one MUST calculate their molar solubilities, s, and then compare them. The larger the value of s, the better (more) soluble the substance is (@ the same T)

molar solubility, s

_ = (K_sp / A^a x B^b)^1/a+b

unrelated

presence of “_” ions in solution (“ionic strength”) changes (typically increases) solubility of insoluble compounds

higher temperature

typically, _ helps to dissolve otherwise insoluble compounds

acids

presence of _ in solution increases solubility of poorly soluble salts of weak acids, e.g. CaCO₃ (CO₃^2- is a conj, base of HCO₃^-)

Le Chatelier’s Principle

When a system at =m is disturbed by a stress (Δ[ ], ΔP, ΔT), the =m shifts in a direction that tends to reduce the effects of that stress

Q < K, Rxn shifts to the right

If products are removed or more reactants are added _

Q > K, Rxn shifts to the left

If products are added or reactants are removed _

Δn_gas ≠ 0

The effects of pressure changes (by means of volume changes) are visible for reactions in which _

P increases

Note: when V decreases, _

Rxn shifts towards fewer moles of gas

If V decreases, P increases, so _

Rxn shifts towards more moles of gas

If V increases, P decreases, so _

gaseous

Warning: You must look at which side of a given reaction has fewer moles of _ species

less sensitive

Reactions which involve liquids and solids are much _ to P (or V) changes

no shift

If pressure changes without volume changes (e.g., by adding another non-reactive gas to the reaction mixture), there is _ to the equilibrium b/c partial pressures remain the same

changes

The =m constant actually _ when temperature changes!

=m shifts to the left, K decreases

For EXOTHERMIC rxns, if T increases _

=m shifts to the right , K increases

For EXOTHERMIC rxns, if T decreases _

=m shifts to the right, K increases

For ENDOTHERMIC rxns, if T increases _

=m shifts to the left, K decreases

For ENDOTHERMIC rxns, if T decreases _

increases the rxn yield

Any change which shifts equilibrium to the right (towards the products) _