AP Exam Review Flashcards: Types of Reactions & Solution Stoichiometry

These flashcards cover key concepts and terms related to types of reactions, solution stoichiometry, gases, thermochemistry, kinetics, equilibrium, acids and bases, applications of thermodynamics, and electrochemistry.

Strong Electrolyte

Substances that completely dissociate into ions in solution; includes strong acids, strong bases, and soluble ionic salts.

Weak Electrolyte

Substances that partially dissociate into ions in solution; typically includes weak acids and weak bases.

Nonelectrolyte

Substances that do not dissociate into ions in solution; includes insoluble ionic salts and covalent molecules.

Net Ionic Equation

An equation that shows only the species that participate in a reaction, omitting spectator ions.

Ideal Gas Law

A relation among the pressure (P), volume (V), temperature (T), and number of moles (n) of a gas, expressed as PV = nRT.

Endothermic Reaction

A reaction that absorbs heat from its surroundings, indicated by a positive change in enthalpy (ΔH).

Exothermic Reaction

A reaction that releases heat to its surroundings, indicated by a negative change in enthalpy (ΔH).

Le Chatelier's Principle

If a system at equilibrium is disturbed, it will shift in a direction that counteracts the disturbance.

Solubility Product Constant (Ksp)

An equilibrium constant that applies to the solubility of sparingly soluble ionic compounds.

Acid-Base Titration

A procedure used to determine the concentration of an acid or base by neutralizing it with a standard solution.

Gibbs Free Energy

A thermodynamic quantity used to predict the spontaneity of a reaction, with negative ΔG indicating a spontaneous reaction.

Redox Reaction

A reaction involving the transfer of electrons, where one species is oxidized (loses electrons) and another is reduced (gains electrons).

What items are all soluble in water

sodium, potassium, ammonium, and nitrate salts

strong acids

HCl, HBr, HI, HClO₄, H₂SO₄, HNO₃

strong bases

Group 1 & 2 hydroxides (Ex: NaOH, Ca(OH)₂, Sr(OH)₂, Ba(OH)₂)

Net Ionic Equation Rules

• Break apart strong acids, strong bases and soluble ionic salts

• Do not break apart weak acids, weak bases, insoluble ionic salts, gases or liquids

when do real gases behave like ideal gases

low pressure and at high temperature

Heat released by the reaction =

heat absorbed by the solution

Heat absorbed by the solution =

mass x increase in temperature x specific heat

The solution is always

 part of the surroundings

we assumed the solution was water with a specific heat of ____  and density of

4.184 J/g°C, 1.0 g/mL

First order

ln [A] vs t

first order graph

straight line

Second order

1/[A] vs t

Second order graph

gives a straight line

The rate of a reactant or product is determined by

the stoichiometry of a reaction

Represent an elementary reaction (one step reaction) as a rate law expression

stoichiometry

Rate laws for reactions that occur in more than one step must be determined by the

method of initial rates and data

The sum of the powers of the reactant concentrations in the rate law is the

overall order of the reaction

the proportionality constant in the rate law is called

the rate constant (k)

 the rate constant (k) is dependent on

temperature

Units of rate constant (k) reflect

the overall reaction order (not always the same)

The rate-determining step in a mechanism is

the slow step and determines the rate law

@ equilibrium, the concentration of reactants and products remain

unchanged

The forward and reverse reactions rates are equal,

results in no observable change to the system

The equilibrium constant expression ONLY includes

gases and aqueous solutions

the equilibrium expression for Kc, Qc

If you reverse a reaction, the K value is

 the reciprocal of the forward reaction. 

(K_reverse = 1/K_forward)

When the balanced equation for a reaction is multiplied by a factor of n

the equilibrium constant for the new reaction is the equilibrium constant for the original reaction raised to the nth power.

(K_new = (K_original)ⁿ

When two reactions with individual K values are combined

the K value for the overall reaction is K₁ x K₂

Reaction quotient (Q)

found the same way as K (law of mass action above) but uses initial concentrations instead of equilibrium concentrations

Q = K

 the system is at equilibrium, no shift occurs

Q < K

the system shifts right to increase the products to reach equilibrium

Q > K

the system shifts left to decrease the products to reach equilibrium

Add a reactant

the reaction shifts right to use up the added reactant

Take away a reactant

the reaction shifts left to make more reactant

Add a product

reaction shifts left to use up the added product

Take away a product

the reaction shifts right to make more product

Increase the volume

decreases pressure. The reaction will shift in the direction to make more moles of gas to increase the pressure.

Decrease the volume

increases pressure. The reaction will shift in the direction to make less moles of gas to decrease the pressure

Common Ion Effect

solubility of a solid is lowered if the solution already contains ions common to the solid.

Q > K_sp

precipitation occurs

Q < K_sp

no precipitation occurs

The value of K_w is temperature dependent, so the pH of pure, neutral water will deviate from 7.0 at

temperatures other than 25°C.

Strong acid + strong base net ionic equation

 H⁺ (aq) + OH^- (aq) H₂O (l)

Weak acid + strong base net ionic equation

HA (aq) + OH^- (aq) A^- (aq) + H₂O (l)

Weak base + strong acid net ionic

B (aq) + H₃O⁺ (aq) HB⁺ (aq) + H₂O (l)

Henderson Hasselbach equation to calculate the pH when you have a concentration of both the weak acid & conjugate base (or weak base and conjugate acid)

pH = pKa + log([base]/[acid])

buffers

resists changes to pH when small amounts of acid or base are added

The BEST buffers have

equal concentrations of the weak acid & conjugate base (or weak base and conjugate acid) AND have a large initial concentration (larger buffer capacity)

If the concentrations are equal w buffer

pH = pKa

Equivalence point - moles of titrant =

moles of analyte originally present

analyte

the substance being measured or analyzed in a titration

titrant

the reagent added in a titration that reacts w/ analyte

Strong Acid + Strong Base, pH

7

Weak Acid + Strong Base, pH

> 7

Weak Base + Strong Acid, pH

<7

Half-way to the Equivalence point (Half-Equivalence point) is only for

titrations of weak acids or weak bases

liquid evaporating to a gas

+ΔS = increase in disorder

water freezing to ice

-ΔS = decrease in disorder

Spontaneous Reaction (thermodynamically favored)

products are favored

ΔG is negative & K > 1

Nonspontaneous Reaction, reactants are favored

ΔG is positive & K < 1

Equilibrium Gibbs Free energy

ΔG = 0 & K = 1

Thermodynamically favored (spontaneous) does not mean a reaction is

fast. Kinetics determines the speed of a reaction.

Loss of electrons is

oxidation

gain electrons is

reduction

Voltaic (Galvanic) Cell

spontaneous redox reaction (E_cell = +)

Concentration Cell

cell compartments are the cell (electrode & solution)

the lower concentration solution is in the anode and the higher concentration is in the cathode.

Electrolytic Cell

1. nonspontaneous redox reaction, must supply a voltage for the cell to operate

Anode

oxidation (mass of the electrode decreases)

Cathode

reduction (mass of the electrode increases)

concentration cell flow

anode to the cathode

concentration cell salt bridge

 connects to anode & cathode and contains ions to maintain a balanced charge

The more positive half reaction STAYS

reduction

Hess law thermochem: Flip the half reaction that should be ___ and change the sign of ___

oxidation, E value

**Never multiply E by an integer even if you have to balance the half reaction!

hess law thermochem: