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86 Cards in this Set

  • Front
  • Back
States of Matter
Solid, Liquid, Gas
Intermolecular Forces required to create a solid: strong or weak?
Intermolecular Forces required for a gas: strong or weak?
The temperature at which liquid and solid are at equilibrium.
Melting Point (or Freezing Point)
0° C/273.15 Kelvin
The temperature at which liquid and gas are at equilibrium.
Boiling Point
100°Celsius/373.15 Kelvin
Enthalpy of Liquid to Solid or Gas to Liquid phase change:
negative (-ΔΤ)
Enthalpy of Solid to Liquid or Liquid to Gas phase change:
positive (+ΔΤ)
Enthalpy of Vaporization
Amount of energy required to completely convert liquid at boiling point to gas at boiling point. There is no temperature change to the substance during vaporization.
Enthalpy of Fusion
Amount of energy required to completely convert a substance from a solid at melting point to a liquid at melting point. There is no temperature change during fusion.
Enthalpy of Fusion of Water
79.7 cal/g
Enthalpy of Vaporization of water
Specific Heat of water
1.00 cal/g°C
Formula for calculating Heat in a reaction:
Heat = mass * ΔΤ * Sp.Heat (cal/g°C)

[ΔT = T2 - T1]
Kinetic Molecular Theory of Gases
-There are no attractive forces between particles of gas
-Space between gas particles is very large
-The average kinetic energy of gas particles is proportional to the Kelvin temperature of the gas.
-Gases experience elastic collisions with each other and the walls of the container
Ideal Gases
gases that obey the kinetic molecular theory
Force per unit area pushing against a surface (P = F/A)
Units of Pressure
Atmosphere (atm)
Millimeters of Mercury (mmHg or Torr)
Pascal (Pa)
Pounds per Square Inch (psi)

1 atm = 760 mmHg/Torr = 14.7 psi = 101,325 Pa
Pressure to Volume
Boyle's Law
Pressue and Volume are inversely proportional (P α 1/V)

PV = k (constant) or P1V1 = P2V2
Volume and Temperature
Charles' Law

Volume is Directly proportional to Temperature (VαT)

V/T = k (constant) or V1/T1 = V2/T2
Pressure and Temperature
Gay-Lusacs Law

Pressure is directly proportional to Temperature (P α T)

P/T = k (constant) or P1/T1 = P2/T2
Combined Gas Law
PV/T = k (constant)


P1V1/T1 = P2V2/T2

Temperature MUST be in Kelvin!
moles and Volume
Avagadro's Law

Volume is directly proportional to moles of Gas (V α n)

V/n = k(constant) or V1/n1 = V2/n2
(n = # of moles of gas)
Ideal Gas Constant (R) using atm
R = 0.0821 L·atm/mol·K
Ideal Gas Constant (R) using mmHg
R = 62.4 L·mmHg/mol·K
Ideal Gas Law
PV = nRT
Dalton's Law of Partial Pressure
the total pressure of a mixture of gases is equal to the sum of the partial pressures of the components of the mixture.

Ptotal = P1 + P2 + P3 + ...

P(total)V = n(total)RT
Standard Temperature and Pressure

0°C (273.15 K) and 1atm (760 mmHg)
Standard Molecular Volume
Volume of one mole of gas at STP

22.4L for ANY gas
-Polar molecules are attracted to one another
-strength of the attraction is approx. 1 kcal/mol (1-2% of covalent bond strength)
-tend to be liquid or solid at room temperature
Hydrogen Bonding
-special kind of dipole-dipole bond
-only present when there is a lone pair of Oxygen, Nitrogen, or Flourine present and a Hydrogen bonded to another O, N or F
-stronger than dipole-dipole (1 kcal/mol)
London Dispersion Forces
-all molecules have this (temporary dipoles)
-results from a temporary polarity caused by the random motion of electrons in a molecule
-STRONGEST force experienced by NON-POLAR molecules, weakest of intermolecular forces
Intermolecular Forces in order or strength
Non-Polar Covalent
Hydrogen Bonding
London Dispersion
Vapor Pressure
The partial pressue of gas molecules in equilibrium with liquid. Increases with temperature.
Relation of Boiling Point to pressure
Directly proportional. BP decreases with lower pressure, increases with higher pressure.
Heterogenous Mixture
A non-uniform mixture with regions of different composition
-murky or opaque
-particles often settle or can be filtered
-ex. orange juice, house paint
Homogenous Mixture
Uniform mixture throughout
-Homogenous solution: particles are small, solution is transparent
-Homogenous colloids: larger particles, murky or opaque
The dissolved substance
The substance (usually liquid) the solute is disolved in.
Solute + Solvent
maximum amount of a substance that will dissolve in a solvent at a given temperature.
How a solid dissolves in a liquid.
Solvation / Hydration
the solvent is water
Like dissolves Like
Liquids with similar intermolecular forces will form solutions. Polar compounds dissolve easily in water. Non-polar compounds dissolve in non-polar liquids.
LeChatlier's Principle
increase in stress on one end of a system in equilibrium will result in a shift to the side of the system with fewer molecules to relieve the stress and restore equilibrium.
Solubility and Pressure
Solubility is Directly proportional to Pressure.

C/P = k (constant) or C1/P1 = C2/P2
(C = solubility)
Formula for calculating Molarity of a solution:
M = moles of solute/Liters of solution
Formula for calculating Weight/Volume %
W/V% = g of solute/mL of solution x 100
Formula for calculating Volume/Volume %
V/V% = volume of solute/volume of solution x 100
(units must be the same!)
Formula for calculating Parts per Million by weight.
PPM = Mass of solute/mass of solution x 10E6
(units must be the same)
Formular for calculating parts per million by volume
PPM = volume of solute/volume of solution x 10E6
(units must be the same)
Formula for Dilution
M1 x V1 = M2 x V2

M1 = original concentration (Molarity)
V1 = initial volume
M2 = desired concentration (Molarity)
V2 = final total volume
V2-V1 = amount of solvent required for dilution
Solid Hydrates
ionic compounds that hold water molecules "trapped" in their lattice.

ex. Epsom Salt, MgSO4 · 7H2O
(each molecule of MgSO4 holds 7 water molecules)
An substance that conducts electricity when dissolved in water.

ionic salts, strong acids & bases

weak acids & bases, non ionic compounds, substances that don't dissolve completely in water

non-soluable ionic compounds
Equivalents of Electrolytes
Units used to describe the amount of ions in bodily fluid.

Eq = molar mass of ions(g)/# of charges on the ion

mEq = 1 Eq/1000
The colligative propertives of a solution depend on the ______ rather than the identity of the solute.
The presence of a solute _______ the vapor pressure of the solution.

Fewer solvent molecules can escape from the solution surface into the gas phase.
The presence of a solute _______ the Boiling Point of the solution.
The presence of a solute ______ the freezing point of a solution.
What determines which solutes will have the greatest effect on the colligative properties of the solution?
Number of particles they dissolve into. The more particles, the higher the HP/lower the MP.
In osmosis, solvent particles always pass from an area of ______ concentration to an area of _______ concentration?
Lower to Greater
The sum of all the molarities in the solution.

0.10 NaCl = 0.20 osmol
What is the auto-ionization of water?
Liquid water dissociates VERY SLIGHTLY into [H30+] and [OH-] ions.

H2O(l) + H2O(l) <---> H3O+(aq) + OH-(aq)

Forward reaction IS NOT favored.
Equilibrium Constant of Water
Kw = [H3O+] x [OH-] = 1.00 x 10E-14
Any substance that produces H3O+ ions when added to water.
How to calculate concentration of an acid from the pH.
H3O+ = 10E-pH
How to calculate pH from the contration of the acid.
pH = -log[H3O+]
Sig Figs for pH
Count the number of sig figs for pH AFTER the decimal point only.

ex. H3O+ - 2.00 x 10E-4 (3 sig figs)
pH = 3.699 (3 sig figs AFTER decimal)
How to calculate [OH-] concentration from [H3O+] concentration
[OH-] = 1.00 x 10e-14/[H30+]
How to calculate [OH-] concentration from pH?
First get [H3O+] concentration:
10e-pH = [H3O+]

Then use equilibrium constant:
[OH-] = 1.00 x 10e-14/[H3O+]
acid/base indicators
--molecules that change color depending on the amount of [H3O+] present

--pH paper, paper coated with universal indicator gives approx pH.

--pH meter, electrodes which measure actual pH
Definition of Acid (Arrhenius)
A substance that provides [H3O+] ions when dissolved in water.
Definition of Base (Arrhenius)
a substance that provides [OH-] ions when dissolved in water.
Definition of Acid (Bronsted-Lowry)
a substance that donates an H+ ion.
Definition of Base (Bronsted-Lowry)
A substance that accepts an H+ ion (donates an electron.) MUST HAVE a lone pair of electrons.
Formula for conjugate acids/bases
B: + HA --> BH+ + :A-

B: is the original base
HA is the original acid

BH+ is the conjugate acid of B:
:A- is the conjugate base of HA
Strong Acids
complete dissociated
Weak Acids
do not dissociate completely, reactions are reversible.
Conjugate bases of Strong acides are _______ bases.
Conjugate bases of weak acides are ______ bases.
The Six Strong Acids
HCIO4 - Perchloric Acid
H2SO4 - Sulfuric Acid
NHO3 - Nitric Acid
HI - Hydroiodic Acid
HBr - Hydrobromic Acid
HCl - Hydrochloric Acid

(all other acids are considered weak)
The Strong Bases
Group I and II cations + OH-
(NaOH, KOH, Mg(OH)2, CA(OH)2)

all other bases are considered weak.
Polyprotic Acids
Acids that have more than one H+ to donate. Each subsequent dissociation will be a weaker reaction.

Ex. H2SO4 is a DIPROTIC acid
H2SO4 + H2O --> H3O+ + HSO4-
HSO4- + H2O <--> H3O+ + SO4-2
Other important Acid/Base reactions:
Acid + metal hydroxide = salt + water

Acid + Carbonate/bicarbonate ion --> carbonic acid --> CO2 + H2O

Acid + NH3 --> ammonium salt(aq)
a procedure for determining the concentration of an acid or base. A known concentration of an acid/base is added to an unknown concentration of a base/acid until pH is neutral. The amount of known substance will equal the concentration of the unknown substance.
Another word for H3O+
Hydronium ion
Another word for OH-
Hydroxyl ion