General Chemistry II Final Exam

Set to study for the final exam

Always Soluble

All group I cations

NH₄⁺

Nitrates (NO₃⁻)

Acetates (CH₃CO₂⁻)

Perchlorates and Chlorates (ClO₄⁻, ClO₃⁻)

Usually Insoluble

Hydroxides (OH⁻)

Phosphates (PO₄³⁻)

Carbonates (CO₃²⁻)

Sulfides (S²⁻)

Hydroxide

OH⁻

Peroxide

O₂²⁻

Cyanide

CN⁻

Carbonate

CO₃²⁻

Hydrogen Carbonate (Bicarbonate)

HCO₃⁻

Acetate

CH₃CO₂⁻

Nitrate

NO₃⁻

Nitrite

NO₂⁻

Ammonium

NH₄⁺

Perchlorate

ClO₄⁻

Chlorate

ClO₃⁻

Chlorite

ClO₂⁻

Hypochlorite

ClO⁻

Sulfate

SO₄²⁻

Sulfite

SO₃²⁻

Hydrogen Sulfate

HSO₄⁻

Hydrogen Sulfite

HSO₃⁻

Permanganate

MnO₄⁻

Dichromate

Cr₂O₇²⁻

Chromate

CrO₄²⁻

Phosphate

PO₄³⁻

Phosphite

PO₃³⁻

Hydrogen Phosphate

HPO₄²⁻

Hydrogen Phosphite

HPO₃²⁻

Dihydrogen Phosphite

H₂PO₃⁻

The Six Strong Acids

Hydrochloric Acid, Hydrobromic Acid, Hydroiodic Acid, Nitric Acid, Perchloric Acid, and Sulfuric Acid

Hydrochloric Acid

HCl (aq)

Hydrobromic Acid

HBr (aq)

Hydroiodic Acid

HI (aq)

Nitric Acid

HNO₃ (aq)

Perchloric Acid

HClO₄ (aq)

Sulfuric Acid

H₂SO₄ (aq)

Chemical Equilibrium

a state of dynamic equilibrium in which the rate of the forward reaction and the rate of the reverse reaction are equal

Law of Mass Action

describes equilibrium condition

Large K Values

the equilibrium favors the products

Small K Values

the equilibrium favors the reactants

LeChatelier's Principle

when stress is applied to a system at equilibrium, the system will respond by relieving the stress

Exothermic Reaction

heat is a product

Endothermic Reaction

heat is a reactant

Common Ion Effect/Salt Effect

an equilibrium will shift if a common ion is added to the solution

Q > K

too many products, favor reverse reaction

Q < K

too many reactants, favor forward reaction

Q = K

reaction is at equilibrium

Acid

proton donor, proton must be bound to a highly electronegative element

Base

proton acceptor

Amphoteric

substance that can react as an acid or a base

Strong Acids

100% dissociation of acidic protons

Weak Acids

less than 100% acid dissociation

Strong Bases

100% base hydrolysis in water

Weak Bases

less than 100% base molecules/ions accept protons

Leveling Effect

limits the acid strength of solutes in a particular solvent

Buffer Solutions

weak acid/weak base conjugate pair that resists pH change

Buffer Capacity

the amount of acid/base that a buffer system can resist before pH starts to change rapidly

Titration Equivalence Point

point where the reaction ended

Titration End Point

point where the indicator changes color

Intramolecular Forces

forces holding a molecule together

Intermolecular Forces

forces holding molecules or aggregates of ions/atoms together

Polar Molecules

asymmetrical molecules that have an electronegativity difference greater than or equal to 0.5

these molecules form dilpoles

The Intermolecular Forces From Strongest to Weakest

Ionic Forces > Hydrogen Bonding > Dipoles > London Forces

London Dispersion Forces

weak temporarily induced dipoles

The Relationship Between Intermolecular Forces and Surface Tension

the stronger the intermolecular forces, the stronger the surface tension

Cohesive Force

attraction within the substance/mixture

Adhesive Force

attraction between different substances

Viscosity

resistance to flow

The Relationship Between Intermolecular Forces and Viscosity

the stronger the intermolecular forces, the higher the viscosity

The Equation for Calculating Heat if Temperature is Changing

Q = mC∆T

The Equation for Calculating Heat When There is a Change in State

Q = m∆H

Solvent

the majority compound

there can only be one in the any solution

Solute

the minority compound

there can be many

Electrolytes

compounds which generate ions when dissolved in water

Strong Electrolyte

100% of the formula units dissolve into ions

Weak Electrolytes

less than 100% of the formula units dissociate into ions

Non-Electrolyte

no ions are formed when dissolved in water

Entropy (S)

the amount of disorder in a system

Soluble Compound

more than 0.02 moles of compound can dissolve in 1.0 L of water

Insoluble Compound

less than 0.02 moles of compound can dissolve in 1.0 L of water

Saturated Solution

solution contains maximum amount of solute possible

Supersaturated Solution

a solution that is carefully prepared with a concentration that exceeds its solubility

Miscibility

liquid-liquid solubility

Miscible Liquids

liquids mix to form a solution

Immiscible Liquids

liquids that do not mix

Colloids

consists of solute particles distributed throughout a solution

Henry’s Law

the amount of gas that can be dissolved in a liquid is directly proportional to the partial pressure of the gas above the liquid

Vapor Pressure

pressure of evaporated molecules above a liquid

Molality

moles of solute per kilogram of solvent

Mole Fraction

the ratio of a solute’s molar amount to the total number of moles of all solution components

Colligative Properties

properties that depend on the concentration of the solute and not the identity of the solute

Raoult’s Law

a solution has a lower vapor pressure than that of a pure solvent

Dalton’s Law

the overall pressure of a gas is the sum of the partial pressures of its components

Osmotic Pressure

the pressure that must be applied to prevent the net movement of water from the solvent to the solution of higher concentration

Hypotonic Solution

external concentration lower than internal concentration

water flows into cell

Hypertonic Solution

external concentration higher than internal concentration

water flows out of cell

Isotonic Solution

external and internal concentrations are equal

Thermodynamics

determines the spontaneity of a chemical reaction

the study of energy, work, and heat

Spontaneous

reaction proceeds in direction written

Nonspontaneous

reaction does not proceed in direction written

Kinetics

determines that rate at which the reaction proceeds

Activation Energy

energy of collision needed to initiate a reaction