Studied by 8 people
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
Note 
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