Chem 105 Test 1 Review

Front 1

States of Matter

Back 1

Solid, Liquid, Gas

Front 2

Intermolecular Forces required to create a solid: strong or weak?

Back 2

Strong

Front 3

Intermolecular Forces required for a gas: strong or weak?

Back 3

weak

Front 4

The temperature at which liquid and solid are at equilibrium.

Back 4

Melting Point (or Freezing Point)
0° C/273.15 Kelvin

Front 5

The temperature at which liquid and gas are at equilibrium.

Back 5

Boiling Point
100°Celsius/373.15 Kelvin

Front 6

Enthalpy of Liquid to Solid or Gas to Liquid phase change:

Back 6

negative (-ΔΤ)

Front 7

Enthalpy of Solid to Liquid or Liquid to Gas phase change:

Back 7

positive (+ΔΤ)

Front 8

Enthalpy of Vaporization

Back 8

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.

Front 9

Enthalpy of Fusion

Back 9

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.

Front 10

Enthalpy of Fusion of Water

Back 10

79.7 cal/g

Front 11

Enthalpy of Vaporization of water

Back 11

???

Front 12

Specific Heat of water

Back 12

1.00 cal/g°C

Front 13

Formula for calculating Heat in a reaction:

Back 13

Heat = mass * ΔΤ * Sp.Heat (cal/g°C)

[ΔT = T2 - T1]

Front 14

Kinetic Molecular Theory of Gases

Back 14

-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

Front 15

Ideal Gases

Back 15

gases that obey the kinetic molecular theory

Front 16

Pressure

Back 16

Force per unit area pushing against a surface (P = F/A)

Front 17

Units of Pressure

Back 17

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

Front 18

Pressure to Volume

Back 18

Boyle's Law
Pressue and Volume are inversely proportional (P α 1/V)

PV = k (constant) or P1V1 = P2V2

Front 19

Volume and Temperature

Back 19

Charles' Law

Volume is Directly proportional to Temperature (VαT)

V/T = k (constant) or V1/T1 = V2/T2

Front 20

Pressure and Temperature

Back 20

Gay-Lusacs Law

Pressure is directly proportional to Temperature (P α T)

P/T = k (constant) or P1/T1 = P2/T2

Front 21

Combined Gas Law

Back 21

PV/T = k (constant)

or

P1V1/T1 = P2V2/T2

Temperature MUST be in Kelvin!

Front 22

moles and Volume

Back 22

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)

Front 23

Ideal Gas Constant (R) using atm

Back 23

R = 0.0821 L·atm/mol·K

Front 24

Ideal Gas Constant (R) using mmHg

Back 24

R = 62.4 L·mmHg/mol·K

Front 25

Ideal Gas Law

Back 25

PV = nRT

Front 26

Dalton's Law of Partial Pressure

Back 26

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

Front 27

STP

Back 27

Standard Temperature and Pressure

0°C (273.15 K) and 1atm (760 mmHg)

Front 28

Standard Molecular Volume

Back 28

Volume of one mole of gas at STP

22.4L for ANY gas

Front 29

Dipole-Dipole

Back 29

-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

Front 30

Hydrogen Bonding

Back 30

-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)

Front 31

London Dispersion Forces

Back 31

-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

Front 32

Intermolecular Forces in order or strength

Back 32

Ionic
Non-Polar Covalent
Hydrogen Bonding
Dipole-Dipole
London Dispersion

Front 33

Vapor Pressure

Back 33

The partial pressue of gas molecules in equilibrium with liquid. Increases with temperature.

Front 34

Relation of Boiling Point to pressure

Back 34

Directly proportional. BP decreases with lower pressure, increases with higher pressure.

Front 35

Heterogenous Mixture

Back 35

A non-uniform mixture with regions of different composition
-murky or opaque
-particles often settle or can be filtered
-ex. orange juice, house paint

Front 36

Homogenous Mixture

Back 36

Uniform mixture throughout
-Homogenous solution: particles are small, solution is transparent
-Homogenous colloids: larger particles, murky or opaque

Front 37

Solute

Back 37

The dissolved substance

Front 38

Solvent

Back 38

The substance (usually liquid) the solute is disolved in.

Front 39

Solution

Back 39

Solute + Solvent

Front 40

Solubility

Back 40

maximum amount of a substance that will dissolve in a solvent at a given temperature.

Front 41

Solvation

Back 41

How a solid dissolves in a liquid.

Front 42

Solvation / Hydration

Back 42

the solvent is water

Front 43

Like dissolves Like

Back 43

Liquids with similar intermolecular forces will form solutions. Polar compounds dissolve easily in water. Non-polar compounds dissolve in non-polar liquids.

Front 44

LeChatlier's Principle

Back 44

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.

Front 45

Solubility and Pressure

Back 45

Solubility is Directly proportional to Pressure.

C/P = k (constant) or C1/P1 = C2/P2
(C = solubility)

Front 46

Formula for calculating Molarity of a solution:

Back 46

M = moles of solute/Liters of solution
(mol/L)

Front 47

Formula for calculating Weight/Volume %

Back 47

W/V% = g of solute/mL of solution x 100

Front 48

Formula for calculating Volume/Volume %

Back 48

V/V% = volume of solute/volume of solution x 100
(units must be the same!)

Front 49

Formula for calculating Parts per Million by weight.

Back 49

PPM = Mass of solute/mass of solution x 10E6
(units must be the same)

Front 50

Formular for calculating parts per million by volume

Back 50

PPM = volume of solute/volume of solution x 10E6
(units must be the same)

Front 51

Formula for Dilution

Back 51

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

Front 52

Solid Hydrates

Back 52

ionic compounds that hold water molecules "trapped" in their lattice.

ex. Epsom Salt, MgSO4 · 7H2O
(each molecule of MgSO4 holds 7 water molecules)

Front 53

Electrolytes

Back 53

An substance that conducts electricity when dissolved in water.

STRONG ELECTROLYTES:
ionic salts, strong acids & bases

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

NON ELECTROLYTES
non-soluable ionic compounds

Front 54

Equivalents of Electrolytes

Back 54

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

Front 55

The colligative propertives of a solution depend on the ______ rather than the identity of the solute.

Back 55

Concentration

Front 56

The presence of a solute _______ the vapor pressure of the solution.

Back 56

Lowers.

Fewer solvent molecules can escape from the solution surface into the gas phase.

Front 57

The presence of a solute _______ the Boiling Point of the solution.

Back 57

RAISES

Front 58

The presence of a solute ______ the freezing point of a solution.

Back 58

Lowers

Front 59

What determines which solutes will have the greatest effect on the colligative properties of the solution?

Back 59

Number of particles they dissolve into. The more particles, the higher the HP/lower the MP.

Front 60

In osmosis, solvent particles always pass from an area of ______ concentration to an area of _______ concentration?

Back 60

Lower to Greater

Front 61

Osmolarity

Back 61

The sum of all the molarities in the solution.

0.10 NaCl = 0.20 osmol

Front 62

What is the auto-ionization of water?

Back 62

Liquid water dissociates VERY SLIGHTLY into [H30+] and [OH-] ions.

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

Forward reaction IS NOT favored.

Front 63

Equilibrium Constant of Water

Back 63

Kw = [H3O+] x [OH-] = 1.00 x 10E-14

Front 64

Acid

Back 64

Any substance that produces H3O+ ions when added to water.

Front 65

How to calculate concentration of an acid from the pH.

Back 65

H3O+ = 10E-pH

Front 66

How to calculate pH from the contration of the acid.

Back 66

pH = -log[H3O+]

Front 67

Sig Figs for pH

Back 67

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)

Front 68

How to calculate [OH-] concentration from [H3O+] concentration

Back 68

[OH-] = 1.00 x 10e-14/[H30+]

Front 69

How to calculate [OH-] concentration from pH?

Back 69

First get [H3O+] concentration:
10e-pH = [H3O+]

Then use equilibrium constant:
[OH-] = 1.00 x 10e-14/[H3O+]

Front 70

acid/base indicators

Back 70

--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

Front 71

Definition of Acid (Arrhenius)

Back 71

A substance that provides [H3O+] ions when dissolved in water.

Front 72

Definition of Base (Arrhenius)

Back 72

a substance that provides [OH-] ions when dissolved in water.

Front 73

Definition of Acid (Bronsted-Lowry)

Back 73

a substance that donates an H+ ion.

Front 74

Definition of Base (Bronsted-Lowry)

Back 74

A substance that accepts an H+ ion (donates an electron.) MUST HAVE a lone pair of electrons.

Front 75

Formula for conjugate acids/bases

Back 75

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

Front 76

Strong Acids

Back 76

complete dissociated

Front 77

Weak Acids

Back 77

do not dissociate completely, reactions are reversible.

Front 78

Conjugate bases of Strong acides are _______ bases.

Back 78

WEAK

Front 79

Conjugate bases of weak acides are ______ bases.

Back 79

Stronger

Front 80

The Six Strong Acids

Back 80

HCIO4 - Perchloric Acid
H2SO4 - Sulfuric Acid
NHO3 - Nitric Acid
HI - Hydroiodic Acid
HBr - Hydrobromic Acid
HCl - Hydrochloric Acid

(all other acids are considered weak)

Front 81

The Strong Bases

Back 81

Group I and II cations + OH-
(NaOH, KOH, Mg(OH)2, CA(OH)2)

all other bases are considered weak.

Front 82

Polyprotic Acids

Back 82

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

Front 83

Other important Acid/Base reactions:

Back 83

Acid + metal hydroxide = salt + water

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

Acid + NH3 --> ammonium salt(aq)

Front 84

Titration

Back 84

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.

Front 85

Another word for H3O+

Back 85

Hydronium ion

Front 86

Another word for OH-

Back 86

Hydroxyl ion