Chem 1040: Test 2

Front 1

Homogeneous vs Heterogeneous equilibrium

Back 1

homo- all same state (gases, aqueous)
hetero- involved rectacts and/or products in a different phase

Front 2

Relationship between Kc and Kp

Back 2

Kp=Kc(RT)^deltan
R= .0821 atmL/molK
T= temp in kelvin (celcius +273.15)
delta n= change in # moles= #mol products-#mol reactants

Front 3

Arrhenius definition of Acid and Base

Back 3

Acid: a substance, that when dissolved in water, increases the concentration of hydrogen ions (H+)

Base: a substance, that when dissolved in water, increases the concentration of hydroxide ions (OH-)

Front 4

Bronsted-Lowry definition of Acid and Base

Back 4

Acid: a proton donor
Base: a proton acceptor

Front 5

What happens when an acid dissolved in water

Back 5

Water acts as a Bronsted-Lowry base and abstracts a proton (H+) from the acid
as a result, a conjugate base of the acid and a hydronum (H3O) ion are formed

Front 6

Strong vs Weak Acids

Back 6

-Strong acids are completely dissociated in water (create incredibly weak bases- almost negligible)
-Weak acids only dissociate partially (conjugate bases are weak bases)

Front 7

How does acid/base strength affect eh side of equilibrium a reaction favors

Back 7

Equilibrium favors the reaction that moves the proton to a stronger base

so in CH3CO2H + H2O => H3O +CH3CO2
the reaction of adding water moves the proton to the strong base, acetate, so equilibrium favors the left side

Front 8

What is amphiprotic and give examples

Back 8

molecule that can be either an acid or base
HCO3-
HSO4-
H2O

Front 9

What is the autoionization of water

Back 9

in pure water a few molecules will dissociate to act as acids or bases forming OH and H30

Front 10

Ion Product Constant of Water

Back 10

equilibrium expression for the autoionization of water
Kc=[H3O+][OH-]
referred to as ion product constant Kw
at 25 C, Kw=1.0 x 10E-14

Front 11

what is pH

Back 11

the negative base-10 logarithm for the concentration of hydronium ion
pH= -log [H3O+]

Front 12

using "p" scales

Back 12

the p in pH stants for negative base-10 log

can use for
pOH = -log[OH-]
pKw= -log Kw

Front 13

using pH to determine acid or base

Back 13

in pure water, pH=7

an acid has a higher [H3O+] then water so pH<7
a base has a lower [H3O+] then water so pH>7

Front 14

How do [H+] and [OH-] compare in acidic and basic solutions

Back 14

neutral: [H+] = [OH-]
acid: [H+] > [OH-]
base: [H+] < [OH-]

Front 15

How do logarithms affect significant figures

Back 15

-when taking the log, use as many digits to the right of the decimal as significant figures in the original
-when inversing, you answer needs as many digits total as there were digits right to the decimal of the log

-log 123 = #.123

Front 16

what is pKw

Back 16

-log (Kw)
-log (1.0 x 10E-14)

pKw=14

Front 17

converting between [OH-] [H+] pOH and pH
* finding pH from [OH]

Back 17

two options
1. find pOH then use pH+pOH=14
2. find [H] by using [H][OH]=1.0 x10E-14 then find pH

Front 18

Characteristivs of a strong acid

Back 18

strong electrolytes
exist totally as ions in aqueous soultions

Front 19

What are the Seven Strong Acids

Back 19

HCl
HBr
HI
HNO3
H2SO4
HClO3
HClO4

Front 20

finding the pH of a monoprotic acid

Back 20

in a monoprotic (one H) acid, the [H+] is the same as the concentration of the acid
so to find the concentration of say .5 M HCl, you assume [H+]=.5

so pH= -log(.5)

Front 21

characteristics of a strong base

Back 21

strong electrolytes
dissociate complely in aqueous solutions

Front 22

what are the strong bases?

Back 22

- alkali metal hydroxides (far left row of periodic table)
-hydroxides of Ca 2+, Sr 2+ and Ba 2+

Front 23

finding the pH of a strong base

Back 23

the pH of a strong base aqueous solution can be determined using the concentration of the strong base
pH of .5 M NaOH = -log (.5)

Front 24

Dissociation Constants

Back 24

Ka and Kb
-the stronger the dissociation constant the stronger the acid or base
Ka= [H3O][A]/[HA]
Kb= [OH][HB]/B

Front 25

Calculating percent ionization

Back 25

percent ionization = [H3O+] at equilibrium/ [Hacid] inital x100

Front 26

relationship between Ka and Kb

Back 26

Ka * Kb= Kw (if you know one you can calculate the other)

Front 27

Reactions of cations and anions with water

Back 27

Anions (negative charge) are bases- will form OH and a conjugate acid
Cations (postive charge) are acids will form H3O and a conjugate base

Front 28

Effect of Cations and Anions on pH

Back 28

-an anion that is the conjugate base of a strong acid will not affect the pH
-an anion that is the conjugate base of a weak acid will increase the pH
-a cation that is the conjugate acid of a weak base will decrease the pH
- cations of the strong arrhenius bases will not affect the pH
-other metal ions will decrease pH

Front 29

Factors affecting the Strength of Acids

Back 29

1. larger the Ka- the stronger the acid
2. in binary acids the more polar and weaker the A-H bond, the more acidic (as you go left and down the periodic table, acidity increases)
3. in oxyacids (with oxygen in them) the more oxygen the more acidic
4.in oxyacids- more acidic as it goes up the periodic table, not down

Front 30

Binary Acids
Oxyacids

Definitions

Back 30

Binary Acids: Basic Formula HA (H is hydrogen and A is other element or ion)

Oxyacids: acid that contains oxygens

Front 31

Definition of Lewis Acid and Base

Back 31

Lewis Acid: an electron pair acceptor *must have empty valence orbital
Lewis Base: an electron pair donor *anything that is a bronsted-lowry base is a lewis base

Front 32

The Common-Ion Effect

Back 32

in a solution where two parts share a common ion (usually H) the H3O may take the concentration of that ion

Front 33

Buffer

Back 33

a solution that resists a change in pH when a small amount of acid or base is added

contains both acid and base so it can neutralize stronger acids and base

Front 34

Henderson-Hasselbalch Equation

Back 34

used to find the pH of a buffer
pH=pKa+log[base]/[acid]
derived from Ka equation
Ka= [H3O][A]/[HA]
Ka= [H3O] *[A]/[HA]
take -log of both sides
pKa=pH-log[base]/[acid]

Front 35

Buffer capactity
Buffer pH range

Back 35

the amount of acid or base a buffer can neutralize before the pH begins to change

the range of pH values oer which a buffer system works

Front 36

Titration

Back 36

a known concentration of base (or acid) is slowly added to a solution of acid (or base)
a pH meter or indicators are used to determine when the solution has reached the equivalence point (stoichiometric amount of acid and base are equal)

Front 37

Factors affecting Solubility

Back 37

-pH: if a substance is a basic anion it will be more soluble in an acidic solution, if a substance is an acidic cation it will be more soluble in a basic solution
-complex ions: metal ions can act as lewis acids and form complex ions with lewis bases in a solvent- this increases solubility
-amphoterism: amphoteric metal oxides and hyrdoxides are soluble in strong acids and base because they can act as either. (amphoteric cations include Al3+, Zn 2+, Sn 2+

Front 38

Will a precipitate form if
Q=Ksp
QQ>Ksp

Back 38

Q=Ksp: at equilibrium and solution is saturated
Q Q>Ksp: the salt will precipitate

Front 39

Fundamental of Acid-Base Equilibria
(4 types of acid-base equilibria)

Back 39

1. Dissociation of a weak acid
any weak acid will dissociate
HA -> H+ + A-
or HA + H2O -> H3O+ + A-

2. Dissociation/ionization of a weak base
any weak base will undergo an ionization
B+H2O -> BH+ + OH-

3. Hydrolysis of anions from weak acids
anions from strong acids do no undergo hydrolysis
anions from weak acids do
A- + H2O -> HA +OH

4. Hydrolysis of cations from weak bases
BH+ -> B + H+