Chem +: Chapter 14, 15, & 18: Equilibrium, Acids, and Bases

Reversible Reactions

Reactions in which products can react and reproduce reactants. Denoted with a two-way arrow.

Equilibrium

State reched where both reactions in a reversible reaction are still occurring, but at the same rate. The rate of the forward reaciton is equal to the rate of the reverse reaction. The concentrations of product and reactant, while not necessarily equal, are constant.

Dynamic Equilibrium

State of equilibirum where reactions are still occurring, but nothing is changing. This occurs in physical systems and chemical systems. An example of this is water evaporation.

LeChatlier’s Principle

If a stress (change) is applied to a system, the system will react to relive the stress (counterbalance that change)

3 potential stresses

1. Change in Concentration


1. No impact for solids/liquids
2. If susbstance is added, reaction will shift to remove that substance (shift to the opposite side)
2. Change in temperature


1. If heat is added, the reaction will shift to remove that heat (depends on exothermic/endothermic)
3. Change in pressure


1. Only impacts gases
2. If pressure is increased, the reaction will shift to decrease that pressure (whichever side of the reaction produces less moles of gas)

Equilibrium Constant (Keq or K)

Ratio of product to reactant concentrations at equilibrium. Constant at constant temperature. Only determined by solution or gas concentration (NOT solid/liquid.) K > 1 → more products. K < 1 → more reactants.

* ( \[C\]^c \* \[D\]^d) / (\[A\]^a \* \[B\]^b)
* \[x\] = concentration of reactant or product
* A and B are reactants, C and D are products
* a, b, c, and d are the coefficients of the reactants and products in the balanced equation
* When there are no gaseous or aqueous products or reactants, a 1 should be put in their place in the formula

Reaction quotient (Q)

Same formula as Q, but it uses the initial concentrations of reactants and products. This can be used to predict how a reaction will shift. If Q > K, the reaction will shift left and produce more reactants. If Q = K, the reaction is already at equilibrium. If Q < K, the reaction will shift and right and produce more products.

Kp

Constant calculated using partial pressures of gaseous reactants and products

Ksp

Describes an equilibirum between a solid ionic compound and a saturated solution of its ions. Determines how much of a slightly soluble or insoluble compound actually dissociates in water.

ICE Method

1. Balanced Equation
2. Write equilibrium expression
3. Determine the **I**nitial concentrations of all elements
4. Determine how the concentrations will **C**hange as the reaction moves toward equilibrium
5. Determine the **E**quilibrium concentrations of all elements
6. Substitute variables into equilibrium expression
7. Solve for x using quadratic formula or the calculator

Properties of acids

Taste sour, react with metals, turn litmus paper red, feel like water, have a higher concentration of hydrogen ions than hydroxide ions, and have a high conductivity in water

Properties of bases

Taste bitter, do not react with metals, turn litmus paper blue, feel slippery, have a higher concentration of hydroxide ions than hydrogen ions, and have a high conductivity in water

Arrhenius concept

Acids donate hydrogen, producing H3O+, a hydronium ion, when dissolved in water. Bases donate hydroxide ions when dissolved in water.

Conjugate Acid

Whichever product gets the hydrogen ion.

Conjugate Base

Whichever product doesn’t get the hydrogen ion

Aphoteric

Substance that can be used as an acid or a base (H2O)

Bronsted-Lowry concept

Same principle for acids as Arrhenius, but recognizes more compounds as bases. They recognize bases as anything that accept hydrogen ions.

Salts

Any ionic compound not containing hydrogen or hydroxide ions. Generally, but not always neutral.

Strong acids and bases

Substances that completely dissociate into ions when dissolved in water.

Strong Acids (must know)

HCl, H2SO4, HClO4, HBr, HNO3, HI, HClO3

Strong Bases (should probably know)

All alkali and alkaline earth hydroxides

pH

Stands for potential of hydrogen ion. The lower this is, the more acidic a substance is. The higher this is, the more basic a substance is. It is equal to -log\[H+\]. Similarly, \[H+\] = 10^-pH.

pOH

Stands for potential of hydroxide ion. The lower this is, the more basic a substance is. The higher this is, the more acidic a substance is. It is equal to -log\[OH-\]. Similarly, \[OH-\] = 10^-pOH.

Relation between pH, pOH, \[H+\], and \[OH-\].

pH + pOH = 14

\[H+\]\[OH-\] = 10^-14

Determining pH and pOH for weak acids and bases.

* Don’t fully dissociate
* Ka and Kb must be used, as well as ICE

Neutralization Reactions

* An acid and a base reacted together produce a salt and water
* 2 methods can be used to determine the amount needed for a perfect neutralization
* Stoichiometry (always works)
* MaVa = MbVb (only works if there is a one-to-one stoichiometric relationship between acid and base)

Titration

Method used to determine unknown concentration of an acid or base

Equivalence Point (Endpoint)

Point in reaction where equal number of moles of acid (H+) and base (OH-) have been added

* pH = 7 when acid and base are both strong
* pH > 7 when acid is weak
* pH < 7 when base is weak

Can be determined using chemical indicators and a pH sensor

Titration curves

Graphs that depict the relation between the volume of an acid or a base and the pH of a substance

Acid/Base properties of Salts

The conjugate base of a weak acid is weak, and the conjugate acid of a weak base is weak. Therefore, when ions form salts, the salts do not have a pH of 7.