Chem 200 Final

Front 1

Like dissolves like means

Back 1

Polar and Ionic substances tend to dissolve in polar solvents
Nonpolar tends to dissolve in nonpoloar

Front 2

Soluble/Not Soluble

1. Hydrogen bonds
2. Dispersion forces
3. Vit A
4. Vit C

Back 2

1. Soluble - stronger IMF
2. Not - weak IMF
3.Soluble in fats (nonpolar)
4.Soluble in H2O

Front 3

Gases in Solution
solubility increases w/mass
because

Back 3

Larger molecules have stronger dispersion IMF

Front 4

Solubility of gas in liquid is directly proportional to

Back 4

It's pressure

Front 5

Henry's Law

Back 5

Sg=kPg

Solubility of gas is equal to the constant for that gas in solvent times the partial pressure of the gas above liquid

Front 6

Generally
As Temp increases, solubility of solid solutes

Back 6

Increases

Front 7

As Temp increases, the solubility of gases

Back 7

decreases (think soda)

Front 8

Solutions:
Saturated

Unsaturated

Super Saturated

Back 8

Solvent holds as much solute as possible at that temp

Less than the maximum amount of solute for that temp

Solvent holds more solute than is normally possible at that temp - unstable, crystallization stimulated by seed crystal/scratch in glass

Front 9

Solutions from solids form by

Back 9

Solvent molecules attracted to surface ions, each ion surrounded by solvent molecules
(solvated)
intermolecular force is ion-dipole

Front 10

Enthalpy of a solution

Three parts

Back 10

Separation of solute particles
- endo
Separation of solvent particles
- endo
Formation of new interactions
- exo

Front 11

Colligative Properties depend only on

Back 11

The number of solute particles present, not the identity of the solute

Front 12

Four colligative properties

Back 12

Vapor pressure lowering
Boiling Point elevation
Melting point depression
Osmotic Pressure

Front 13

Colligative properties of VP due to

Back 13

As solute molecules are added to a solution, the solvent becomes less volatile (= decreased vapor pressure)
solute - solvent interaction

Front 14

BP Elevation equation

Back 14

Change in Tb = Kb* m (molality)
Then: add to BPi

Front 15

FP depression equation

Back 15

Change in Tf=Kf * m (molality)
Then subtract from FPi

Front 16

Electrolytes show greater changes because regarding colligative properties because

Back 16

The number of particles is greater when they dissociate
ie: NaCl dissociates into two ions

Front 17

Osmosis

Back 17

Semi-permeable membranes allow some particles to pass through while blocking others

Front 18

Osmosis - net movement of H2O

Back 18

from the area of higher solvent concentration (lower solute) to the area of lower solvent concentration (higher solute)

Front 19

Osmotic Pressure

Back 19

The pressure required to stop osmosis
pi is Osmosis in blood cells

Front 20

Hypertonic

Back 20

The solute concentration outside of cell is greater than inside.
Water flows out
Crenation results

Front 21

Hypotonic

Back 21

Solute concentration is greater inside the cell than out
Water flows in
Cell lyses

Front 22

Ideal Gas Law

Back 22

PV=nRT

Front 23

Partial Pressure equation

Back 23

P1/P2 =M1/M2

Pressure of gas 1/gas 2 is equal to the moles of gas1/gas2

Front 24

Rate of gas 1 compared to rate of gas 2

Back 24

r1/r2 = square root of M1/M2 so:
Molar mass of 1 times (r1/r2)squared =Molar mass of 2

Front 25

Binary Molecular Compounds
covalent
non-metal-non-metal

naming

Back 25

name the number of atoms of the first, then the second element using Greek Prefixes
add "ide to the end
*excluding mono on first element

Front 26

Greek prefixes through 10

Back 26

mono hexa
di hepta
tri octa
tetra nona
penta deca

Front 27

Except for groups 1a and 2a

4 metal ions with one possible charge

Back 27

Ag+
Zn2+
Cd2+
Al3+

Front 28

Variable Charge
1. Iron

2.Copper

Back 28

Fe (II) and Fe(III)

Cu(+) and Cu(II)

Front 29

Variable Charge
3.Lead

4. Tin

Back 29

PB(II) and Pb(IV)

Sn(II) and Sn (IV)

Front 30

Variable Charge

5. Mercury

Back 30

Hg(II) = mercury (II)

Hg2 (2+) = mercury 1

Front 31

Name Binary Acids

Back 31

Hydro....ic
ex:
Hydrochloric
Hydrobromic

Front 32

Nam Oxy Acids

Back 32

Contain polyatomic ions
#of H = the charge on the ion
anions ending in ate = ic
anions ending in ite = ous

Front 33

Hydrates

Back 33

ionic compounds with an exact number of H2O molecules bound to each crystal unit

Front 34

Naming Hydrates

Back 34

use Greek prefixes to indicate the number of H2O's
ex. MgSo4 * 7H2O
- magnesium sulfate heptahydrate

Front 35

Electrolyte

Back 35

Forms ions in a solution
conducts electricity

Front 36

Electrolyte examples

Back 36

Strong acids
Strong bases
Soluble Salts

Front 37

Strong Acids

Back 37

HCl
HNO3 - nitric acid
H2SO4 - sulfuric acid
HB - hydrobromic
HI - hydroionic
HClO4 - perchloric acid
HClO3 - chloric acid

Front 38

Strong Bases

Back 38

LiOH
NaOH
KOH
Ca(OH)2
Sr(OH)2
Ba(OH)2

Front 39

Sol rules:

Always soluble

Back 39

Alkali metals - group 1a
NH4
NO3,C2H3O2

Front 40

Cl-,Br-, I- soluble except with...

Back 40

Ag+, Hg2+, Pb2+

Front 41

OH- and
S- are only soluble with

Back 41

Sr, Ba, Ca, and the alkali metals
and NH4

Front 42

SO4 2- is soluble except

Back 42

Sr2+, Ba2+, Hg2 2+, and Pb2+

Front 43

CO3 2- and PO4 3-
are only soluble with

Back 43

NH4 and the alkali metals

Front 44

Oxidation numbers:

Atoms in elemental form

Back 44

are always zero

Front 45

Ox #'s

monatomic ions

Back 45

the ox# equals the charge

Front 46

Ox#'s

hydrogen

Back 46

+1 when bonded to nonmetals
-1 when bonded to metals

Front 47

Ox#'s

Flourine
other halogens

Back 47

-1 in all compounds

-1 in most binary compounds

Front 48

oxidation numbers increase when

Back 48

an electron is lost

(oxidation)

Front 49

oxidation numbers decrease when

Back 49

an electron is gained

(reduction)

Front 50

The oxidizing agent is reduced and

Back 50

The reducing agent is oxidized