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FINAL Flashcards

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AP® Chemistry flashcards Periodic Table flashcards
1
Q

electronegativity

A

an atom’s ability to attract another atom’s electrons

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2
Q

electronegativity trend

A

increases up (smaller distance to the nucleus) and to the right (smaller radius, more protons)

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3
Q

nonpolar meaning and number

A

equal share of electrons, less than 0.5

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4
Q

polar meaning and number

A

unequal share of electrons, between 0.5 and 1.7

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5
Q

ionic meaning and number

A

transferred electrons, greater than 1.7

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6
Q

metallic character increases…

A

diagonally down

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7
Q

when a bond is broken energy is…

A

absorbed

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8
Q

when a bond is formed energy is…

A

released

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9
Q

lattice energy

A

energy needed to separate ions in ionic compounds, combining ions will release energy, represented w/ coulombs law

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10
Q

properties of ionic compounds

A

form crystals, nonvolatile (slow to vaporize), soluble in polar solvents, high melting and boiling points, hard, brittle, conduct electricity when melted, good insulators when solid, strong

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11
Q

polarity of the periodic table

A

least polar = closest together, most polar = farthest apart

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12
Q

partial negative and positive charges

A

closer to F = partial negative, farther from F = partial positive

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13
Q

colombic attraction in metals

A

number of valence ele determines dumber of electrons in sea, charge of cations and number of ele increase the attractions are greater, when radius decreases so does attraction

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14
Q

properties of metals

A

lustrous, malleable, ductile, good conductors of electricity, high BP and MP, strong

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15
Q

interstitial alloys

A

fit in between metal atoms holes

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16
Q

substitutional alloys

A

atoms have similar radii, replace atoms in the lattice

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17
Q

the ____ electronegative atoms go in the center of a lds

A

least

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18
Q

formal charge

A

the charge an atom would have if all of the electrons in a covalent bond were shared equally

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19
Q

how to calculate formal charge

A

valence electrons - number of lines and dots

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20
Q

best formal charges

A

add up to charge, if neutral all must be zero, central atom has a charge of zero, more electronegative is negative, less electronegative is positive or zero, low

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21
Q

fewer than 8 electrons

A

Li, Be, B

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22
Q

more than 8 electrons

A

elements in period 3 or below

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23
Q

localized electrons

A

specifically at one atom or shared between 2

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24
Q

delocalized electrons

A

shared by multiple atoms

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25
Q

drawing organic molecules

A

connect c, evenly distribute Hs, draw order given

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26
Q

hybrid orbitals

A

domain counters, when 2 ele domains mix they created 1 less hybridization

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27
Q

sigma bonds

A

first bond between two atoms head to head (single, never more than 1)

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28
Q

pi bonds

A

side to side overlap

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29
Q

bond order

A

the number of bonds shared/the number of atoms

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30
Q

polar molecules

A

have a dipole moment (difference in electronegativity), lone pairs

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31
Q

metals are both malleable and ductile. The best explanation for these properties is that the electrons involved in bonding among metal atoms are

A

equally shared and form nondirectional bonds

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32
Q

Atomic symbol

A

Top # mass (protons + neutrons)
Bottom # atomic number (protons)

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33
Q

majority of mass

A

pro and neu

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34
Q

Shielding effect

A

Core electrons shield the nucleus from the valence electrons

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35
Q

Effective nuclear charge

A

the positive charge experienced by valence electrons
zeff = atomic number - number of core electrons
–> number of valence electrons

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36
Q

ionization energy trend

A

when going across there are more electrons bound to the nucleus because the force increases because the charge increases, as you move down a family the atomic radius gets bigger and the protons and electrons spread out making less energy needed to separate them

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37
Q

Radius trend

A

Going right electrons get pulled tighter
Moving down gains energy levels

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38
Q

Electron affinity

A

Change in energy that occurs when an electron is added

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39
Q

Electron affinity trend

A

Inc up and right as the distance between electrons and nucleus decreases with more electrons and added protons make it easier to gain an electron

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40
Q

Electron affinity outliers

A

Period 2 is low (full shell)
Group 15 is low

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41
Q

Charges to memorize

A

Ag+
Cd2+
Zn2+

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42
Q

Empirical formula vs molecular

A

Simple vs actual

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43
Q

Periodic trends in general

A

Across = More protons, more valence electrons
Down = Greater distance from nucleus

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44
Q

are molecules with resonance polar or nonpolar

A

nonpolar
Equal contributions of the three resonance forms to the structure of the ion result in a planar ion with three equivalent polar nitrogen-oxygen bonds. Due to the trigonal planar symmetry of the
ion, the negative dipoles cancel out.

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45
Q

solubility

A

like dissolves like

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46
Q

if the same bonds, the atom with more ____ has the largest bond dissociation energy…

A

protons

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47
Q

preferred LDS

A

negative formal charge on most electronegative, central has zero

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48
Q

dipole-dipole

A

polar to polar
more polar, greater attraction
permanent dipole

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49
Q

LDFs strength

A

weak
more surface area –> stronger

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50
Q

only IMF in nonpolar molecules

A

LDFs

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51
Q

network covalent

A

SiO2, SiC, C (graphite and diamonds)

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51
Q

dipole-dipole bp

A

~equal mass and size, more polar, higher bp

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52
Q

viscosity increases

A

with IMF strength

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53
Q

capillary action w/ strength

A

better capillary, stronger IMFs

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54
Q

capillary action

A

rise of liq up tube

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55
Q

ionic solids

A

metal nonmetal
lattice composed of alternately charged ions
small ions (ions w higher charges) will have stronger attractions
think salt for characteristics

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56
Q

metallic solids

A

metal metal
sea
mg ribbon

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57
Q

covalent-network solids

A

joined by extensive network of covalent bonds in pattern
diamonds

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58
Q

molecular solids

A

held by weak forces
nonmetal nonmetal
think water
don’t conduct

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59
Q

gas law

A

p1v1/t1n1=p2v2/t2n2

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60
Q

real gasses do not behave ideally when…

A

at high pressures and low temperatures
volume of particles and intermolecular attractions are now significant

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61
Q

strength of dispersion forces tends to

A

increase with molecular weight

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62
Q

miscibility

A

ability of two substances to mix without separating
similar materials are miscible in each other

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63
Q

polar solvents dissolve

A

polar and some ionic solutes
polar- molecules spread out with broken imfs
ionic- dissociate

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64
Q

in aqueous solutions

A

water molecules surround ions to dissolve them

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65
Q

paper chromatography

A

more similar in polarity sample is to solvent, farther it travels

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66
Q

thin layer chromatography

A

more polar, shorter distance
used with samples not colored to naked eye - uv (amino acids)

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67
Q

column (method) chromatography

A

steel wool with gel in buret, mixture placed at top, flushed with nonpolar solvent
used more for separation than analysis
more polar, slower

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68
Q

higher frequency

A

high energy

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69
Q

spectroscopy

A

study how matter interacts with radiation

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70
Q

microwave spectroscopy

A

rotational
lower in energy than visible light

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71
Q

infrared spectroscopy

A

vibrational
higher in energy than microwaves, lower than visible light
lighter and stronger bonds, higher vibrational frequency

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72
Q

ultraviolet (visible light) spectroscopy

A

transitions in energy levels
absorbing – high energy
measure light absorbed, determine concentration

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73
Q

photoelectric effect

A

when light shines on surface of a metal, electrons can be ejected from surface

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74
Q

best choice for wavelength on a graph

A

where absorbance close to 1

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75
Q

beer-lambert law

A

on equation sheet
absorbance = slope * path length (1) * concentration (M)

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76
Q

soluble in water…

A

all sodium (Na), potassium (K), ammonium (NH4), nitrate (NO3) salts

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77
Q

if imfs break strength rules

A

due to molar mass, more opporunities

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78
Q

in order to dissolve

A

need strong attraction

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79
Q

pv

A

indirect

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80
Q

pt

A

direct

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81
Q

tv

A

direct

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82
Q

vn

A

direct

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83
Q

do all gas molecules have the same speed?

A

No! Based on kinetic molecular theory, gas particles have a distribution of speeds that is dependent on the temperature

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84
Q

molar absorbity graph outliers

A

too high ~ contamination from solution of higher M
too low ~ contaminations from water of solution of lower M

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85
Q

if given ke dist you can find

A

molecular mass of each
molecular masses are from high to low (same as peaks)

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86
Q

h20 is polar or nonpolar

A

polar

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87
Q

if a solution does not conduct electricity

A

ionization does not occur

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88
Q

absorption with fingerprints on cuvette

A

absorption curve higher bc more light absorbed
fingerprints scatter light, less light would pass and reach detector so more light would be thought to be absorbed

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89
Q

if at STP all vessels have

A

same number of molecules (1 mol or A)

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90
Q

for pressure questions use

A

p1v1/n1t1 if some is help constant and only plug in what changes

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91
Q

make sure wavelength is in

A

meters

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92
Q

mixtures

A

can be separated physically

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93
Q

pure substances

A

can be separated chemically

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94
Q

physical changes

A

changes in matter that do not change the chemical composition of a substance
change in properties
imfs are broken/formed

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95
Q

physical changes examples

A

phase changes, temperature, volume, mixture separation, dissolving

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96
Q

chemical changes

A

result in new substances because the chemical composition has changed
bonds are broken/formed

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97
Q

examples of chemical changes

A

oxidation and decomposition

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98
Q

signs of a chemical change

A

(unexpected) color change, production of odor, change in temp (w/o adding or removing energy), evolution of a gas (bubbles), precipitate

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99
Q

single replacement reaction

A

occur when an element reacts with a compound producing a different element and a different compound
A + BC –> AC + B
redox

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100
Q

to determine which substance is replacing in SR

A

look at signs, must be the same

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101
Q

double replacement aka

A

precipitation/neutralization

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102
Q

combustion

A

CX + O2 –> CO2 + H2O

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103
Q

different types of dissolving

A

ionic compounds - dissociation
molecular compounds - interact/react with water

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104
Q

solubility of ionic compounds

A

Na, K, Ammonium (NH4), Nitrate (NO3) salts

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105
Q

all nonpolar substances are

A

not soluble

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106
Q

if a substance is not soluble it will

A

form a precipitate

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107
Q

precipitation reaction

A

two solutions containing soluble slats are mixed sometimes an insoluble salt will be produced

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108
Q

molecular equation

A

complete chemical equation

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109
Q

ionic equation

A

strong acids, bases, and soluble salts are broken into ions with charges

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110
Q

net ionic equations

A

ionic but get rid of spectator ions

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111
Q

things included in particle diagram

A

state of matter
amount
species

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112
Q

aqueous solution in a particle diagram

A

ions separated, water with dipoles in shape
–if not ionic interactions between water and solute dipoles

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113
Q

species in particle diagram

A

make a key
show charges (if in H2O)
appropriate sizing

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114
Q

percent yield

A

actual/theoretical

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115
Q

titration setup

A
  1. wash buret and tools with water
  2. rinse with acid/base
  3. fill with liquid and record the starting value
  4. titrate it
  5. read final mark and subtract to get the total volume
  6. repeat
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116
Q

if double replacement…

A

check if h donated/gained
–> acid base
check if aq, aq, aq, s
–> precipitaion

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117
Q

if not DR…

A

check oxidation numbers! if changed
–> redox

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118
Q

nonmetals tend to have ___ oxidation numbers when in compounds

A

negative

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119
Q

oxidation number of oxygen if peroxide

A

-1

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120
Q

the more electronegative nonmetal will…

A

get the negative oxidation number

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121
Q

oxidation number of halogens

A

-1

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122
Q

ionic compounds charge order

A

positive than negative

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123
Q

rate of a reaction is measured using

A

the concentration of a reaction over change in time

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124
Q

rate of reactant vs. product (if 1:1)

A

equal and opposite

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125
Q

rate of reactant vs. product (if NOT 1:1)

A

-1/a (rate) = 1/c (rate)
where a and c are the leading coefficients

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126
Q

in order to react, molecules must react…

A

with a sufficient amount of energy and must be in the right orientation
collision theory

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127
Q

factors that affect reaction rate

A

physical state/surface area of reactants
reactant concentrations
pressure
temperature
catalysts

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128
Q

physical state/surface area of reactants affect on reaction rate

A

homogeneous reactions are often faster
heterogeneous reactions that involve solids are faster if the surface area is increased
–increased SA increases the number of particles that are exposed to react and leads to a greater frequency of collisions

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129
Q

reactant concentrations
affect on reaction rate

A

increasing reactant concentration generally increases reaction rate
–since there are more molecules, more collisions occur

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130
Q

catalysts affect

A

the kinds of collisions

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131
Q

rate order

A

tells us how the concentration of a reactant specifically affects the rate of the reaction

132
Q

to find order

A

see how the rate changes by change in concentration to the power of the order number

133
Q

when finding order make sure…

A

one reactant has a constant concentration

134
Q

rate law

A

tells us how each reactant in the chemical equation affects the rate of the reaction
rate = k[A]^x[B]^y… where x and y are the orders

135
Q

to write a rate law..

A

find orders of concentrations
plug in an experiment
solve for k
plug in new numbers using k

136
Q

ratio method

A

when we cannot visually determine the order, we can set up ratios of rate laws for different trials
rate 1/rate 2= rate law/rate law

137
Q

differential rate laws

A

rate dependent on concentration only
only able to determine initial rate

138
Q

integrated rate laws

A

if we want to find the rate or concentration at any point in time
–use a linear graph!

139
Q

radioactive oder

A

first

140
Q

half life

A

amount of time it takes for one half of a reactant to be used up

141
Q

use of integrated rate laws

A

linear = winner

142
Q

transition state on an energy vs time graph

A

peak of hill

143
Q

breaking bonds is

A

endothermic

144
Q

forming bonds is

A

exothermic

145
Q

largest activation energy =

A

slowest step

146
Q

intermediates on energy vs time

A

valleys of catalyst

147
Q

number of peaks in diagram

A

number of steps to reaction

148
Q

catalysts

A

change the path and lower the activation energy
get them back after a reaction and can reuse them over and over again
will appear as a reactant in rate determining step and a product of a later step

149
Q

intermediate

A

species that are neither reactants nor products, they are formed during the reaction and then consumed by the reaction

150
Q

why more steps is common

A

more likely less species are able to collide w correct orientation and sufficient energy

151
Q

equilibrium

A

when rate of forwards reaction is equal to the rate of the reverse reaction
k values are equal and opposite

152
Q

overall order

A

sum of individual orders

153
Q

slope of rate law graphs (if liner) =

A

k

154
Q

rate law of slow step =

A

rate law of overall equation

155
Q

if slow/rate determining step isn’t first

A

must use all reactants up to slow step (w/o intermidiates)

156
Q

highest hump on multi-step reaction

A

slow step

157
Q

rate of disappearance

A

finial concentration-initial concentration/time

158
Q

-Q warm object/-MCT warm object =

A

+Q cooler object, +MCT cooler object

159
Q

Hvap=

A

-Hcond

160
Q

omit from eq expression

A

all solids and pure liquids (anything labeled l)

161
Q

only way to change equilibrium constant

A

changing temperature

162
Q

equilibrium constant unit

A

k has no units

163
Q

adding equations and altering k

A

flip equation, take reciprocal of k
double equation, raise k ^2
multiply k values together

164
Q

dissolving solute-solvent

A

exothermic
don’t need to know why just know more IMFs formed than broken

165
Q

dissolving solute-solute

A

endothermic
don’t need to know why just know more IMFs broken than formed

166
Q

dissolving solvent-solvent

A

endothermic
don’t need to know why just know more IMFs broken than formed

167
Q

heat transfer

A

hot to cold

168
Q

thermal equilibrium

A

when two substances result in the same temperature after a heat transfer
—they have the same average KE

169
Q

molar heat capacity

A

specific heat capacity * molar mass

170
Q

we use constant-pressure colorimetry when…

A

observing heat transfers
observing reactions between solutions

171
Q

hess’s law

A

delta h for overall reaction is sum of enthalpy changes for individual steps

172
Q

going up a phase change graph the reaction is

A

endothermic

173
Q

why would something level out at a constant temp?

A

the molecules will have the same average kinetic energy. . The temperature is no longer changing at this point on the graph because thermal equilibrium was reached.

174
Q

when to use bonds broken-bonds formed vs hproducts-hreactants

A

based on info given
1st one if given bond enthalpies second if given enthalpies of formation

175
Q

constant-volume bomb calorimetry

A

filed with O2 gas before placed in a bucket
sample ignited electrically
heat produced can be determined by measuring temp increase in the surrounding water
often used to determine caloric value of food
–more specific/accurate
–used for precise measurements

176
Q

K (rate law) unit

A

1/M^(order-1)s

177
Q

K (equilibrium) unit

A

none!

178
Q

path length for beer-lambert law

A

not the same as wavelength!! usually given wavelength in nm which is just the most effective wavelength for the specific substance, path length is usually 1cm, look at given units to determine unit

179
Q

specific heat of liquid water

A

4.18J/g*C

180
Q

1mL water =

A

1 g water

181
Q

q = delta h when

A

constant pressure
constant volume

182
Q

in a combustion rxn states of matter

A

assume all gas unless stated otherwise

183
Q

q=

A

mcdeltat (change in temp)
n(amount based on units)*vap/fus

184
Q

if the forward reaction is exothermic…

A

the reaction will be favored (smaller Ea)
there will be more products in equilibrium mixture
opposite is true for endothermic

185
Q

equilibrium constant with molarity

A

[products]^coefficient/[reactants]^coefficient

185
Q

homogenous equilibrium

A

all reactants and products in the same state

185
Q

favored direction

A

if favored products –>
if favored reactants <–

186
Q

change in volume or pressure

A

higher volume/lower pressure favors the side of the equation with more moles (and vice versa)

186
Q

eliminated from eq expression

A

solids pure liquids

187
Q

bigger solubility number

A

more soluble

188
Q

common ion effect

A

if present solubility of the salt will decrease (less room)

189
Q

unsaturated

A

hasn’t reached equilibrium yet
not everyting dissolved

190
Q

will a precipitate form

A

Q<K>K precipitate</K>

191
Q

strong acids

A

HI
HCl
HBr
H2SO4
HNO3
HClO4
HClO3

192
Q

strong bases

A

group 1 and 2 hydroxides

193
Q

weak acids

A

equilibrium

194
Q

percent dissociation

A

[H+]/[acidi]

195
Q

equivalence point graph

A

point of inflection

196
Q

halfway point graph

A

halfway to inflection
pH=pKa

197
Q

if both N and H probs

A

weak base

198
Q

choose indicator that has pKa close to

A

pH of equivalence point

199
Q

buffer

A

wa and cb

200
Q

pH of buffer

A

pH=pKa+log[CB]/[WA]

201
Q

volume and entropy

A

more volume, more entropy

202
Q

k and thermodynamic favorability

A

k>1 = favored
k<1= not favored

203
Q

cat or an increase in mass

A

fat cat

204
Q

fill salt bridge with

A

ions that won’t react

205
Q

use nerst for

A

nonstandard conditions

206
Q

nerst

A

E=E-(RT/nF)lnQ

207
Q

size of Q in nerst

A

+Q~ lowers voltage
-Q~ higher voltage

208
Q

base

A

a negative ion that can accept an H+ ion in solution or has at least one lone pair of electrons
a substance that donates OH- ions to a solution

209
Q

if [H+] > [OH-]

A

the solution is acidic

210
Q

if [H+] < [OH-]

A

the solution is basic

211
Q

strong acids pH

A

close to 0

212
Q

weak acids pH

A

close to 7

213
Q

strong bases pH

A

close to 14

214
Q

weak bases pH

A

close to 7

215
Q

[H+] = (strong acid)

A

[acid]

216
Q

Ka =

A

for weak acids
equation sheet, when HA <–> H + A-

217
Q

the larger the Ka,

A

the stronger the (weak) acid

218
Q

percent ionization

A

measurement of the extent of ionization (usually of an acid)
-a strong acid will have 100%
–a weak acid will have a much lower %

219
Q

weak bases

A

equation: B + H2O <—-> HB+ + OH-

220
Q

weak bases have…

A

lone pairs or a negative charge

221
Q

pH depends on

A

what is remaining in excess

222
Q

anytime ONE reactant is strong, the reaction will

A

go to completion

223
Q

an acid’s/base’s strength is determined by…

A

it’s conjugate

224
Q

a strong acid/base will create…

A

a non-reactive conjugate

225
Q

a non-reactive substance

A

does not react
does not have the characteristics of an acid or a base

226
Q

in every acid-base reaction, equilibrium favors transfer of the H+ ion from

A

the stronger acid to the stronger base

227
Q

organic molecules with more oxygen atoms will be

A

more acidic/stronger

228
Q

to determine if a salt is acidic, basic, or neutral

A

split the salt into its ions
see if the cation and anion will attract H+ or OH-
see if the new compounds are strong or weak
-if both strong/weak = neutral
-if SA but WB = acidic
if SB but WA = basic

229
Q

if a substance has a basic anion (-), the solid will be…

A

more soluble in an acidic solution

230
Q

solubility increases when

A

able to react and form a new product in the liquid

231
Q

third law of thermodynamics

A

the entropy of a pure crystalline substance at absolute zero is 0
–in a perfect lattice there is only one microstate

232
Q

kinetic control

A

the idea that sometimes spontaneous reactions may not occur
—determines if a thermodynamically favorable reaction will actually occur

233
Q

some reasons reactions with a negative delta G do not occur

A

-High activation energy
-Unfavorable orientations
–high numbers of particles that must collide simultaneously

234
Q

free energy at eq

A

delta g = -RTlnK

235
Q

Size of K and thermodynamic favorability

A

When K>1, the term lnK is positive and G is negative, thus the reaction is thermodynamically favorable.

236
Q

concentration cell vs. voltaic cell

A

concentration - same ion for electrodes, different concentrations

237
Q

concentration cells

A

oxidation occurs at lower molarity, electrons move to solution with more positive ions

238
Q

best ionization energy

A

smallest

239
Q

electronegativity of horizontal neighbors

A

nonpolar

240
Q

CH is…

A

nonpolar

241
Q

an element with itself is…

A

nonpolar

242
Q

H fits in to all trends in its location except…

A

electronegativity, close to C

243
Q

if K«1, concentrations and favorability

A

products &laquo_space;reactants
small K indicates products not favored, doesn’t[t indicate reactants are
equilibrium is established when hardly any products have been formed (consumption of reactants is not favored).

244
Q

1 M HCl = 100% dissociated, why is 0.01 M HCN <1% dissociated

A

eq constant of HCN is smaller than HCl

245
Q

molar solubility equation

A

Ksp = [A+]^n [B-]^m

246
Q

molar solubility

A

the amount that is actually dissolved

247
Q

molar solubility of reactant =

A

molar solubility of products (if 1:1)
– if not 1:1 use mole ratio

248
Q

if Q<K

A

rate of forward rxn > rate of reverse rxn

249
Q

sign equilibrium has occurred

A

constant P, not constant mass, law of conservation of mass means mass is always constant

250
Q

large Kp value

A

reaction will proceed to almost equilibrium, but because of excess reactants some reactants are left (the limiting reactant)

251
Q

if Q<K

A

forward rxn is favored

252
Q

as pH decreases [H+]

A

increases

253
Q

if temp increased as product rxn shifts to

A

reactants

254
Q

can ignore. if

A

k«1

255
Q

when does a precipitate form

A

when Q>K

256
Q

SA

A

SA –> H+ + ion-
ion- is nonreactice

257
Q

[SA]=

A

[H+]

258
Q

SB

A

SB –> OH- + ion+
ion+ is nonreactive

259
Q

[SB] =

A

[OH-]

260
Q

WA

A

WA + H2O <–> H3O+ + ion-
ion- is reactive

261
Q

[WA]

A

use ICE table to solve

262
Q

WB

A

SB + H2O <–> OH- + ion+
ion+ is reactive

263
Q

[WB]

A

use ICE table to solve

264
Q

SA + SA

A

not a reaction
find total moles of H+ to get [H+]
–moles of H+ is each acid/total volume = [H+]

265
Q

SB + SB

A

not a reaction
find total moles of OH- to get [OH-]
–moles of OH- is each base/total volume = [OH-]

266
Q

SA + SB

A

SA + SB –> H2O + salt
salt is non reactive

267
Q

pH of SA + SB

A

BCA table to solve for moles of excess
-if SA excess, [H+]=moles of left over H+/total volume
-if SB excess, [OH-]=moles of left over OH-/total volume
-if neither in excess [H+]=[OH-], pH=7

268
Q

SA+WB

A

SA + WB –> CB + CA
CB of a SA is nonreactive
CA of a WB is reactive

269
Q

pH of SA + WB

A

BCA table to solve for moles of excess
-if SA excess, CB is nonreactive, [H+]=moles H+ left over/total volume
-if WB excess, CA is reactive and creates a buffer
–pOH=pKb+log([CA]/[WB])
-if neither is excess, second BCA table with reactive CA in water to solve for moles of H+

270
Q

SB + WA

A

SB + WA –> CA + CB
CB of WA is reactive
CA of SB is non reactive

271
Q

SB + WA pH

A

BCA table to solve for moles of excess
-if SB excess, CA is nonreactive, [OH-]=moles OH- left over/total volume
-if WA excess, CB is reactive and creates a buffer
–pH=pKa+log([CB]/[WA])
-if neither is excess, second BCA table with reactive CB in water to solve for moles of OH-

272
Q

Acid + Base –>

A

Salt + Water

273
Q

M in M1V1

A

molarity

274
Q

increasing concentration of buffer components while keeping ration constant

A

keeps PH same but increases capacity of buffer

274
Q

when a buffer has more CA than B

A

it has a greater capacity for added base than acid and vice versa

275
Q

in pure water pH=

A

pOH, regardless of temperature

276
Q

acid strength decreases as

A

electronegativity decreases

277
Q

percent ionization formula

A

[H+]eq/[HA]initial

278
Q

pKa+pKb=

A

14

279
Q

[H3O+]=[OH-]

A

ALWAYS in pure water

280
Q

initial concentration of base decreases

A

smaller amount of OH- is produced
–higher pOH
–lower pH

281
Q

initial concentration of acid increases

A

more H3O+ releases
–decreased PH

282
Q

half eq point of WA + SB

A

[HA]=[A-]

283
Q

Ka values strength

A

bigger Ka = stronger
-more electronegative

284
Q

the more stable a CB is the

A

stronger the WA is

285
Q

weaker base

A

large electronegativity

286
Q

if more CB than A, solution has greater buffer capacity for

A

acid

287
Q

k>1

A

reaction favors products, exothermic
reactants to products

288
Q

k<1

A

reaction favors reactants, endothermic
products to reactants

289
Q

Q vs. K

A

nature wants Q=K(eq)
if Q<K>K proceed to reactants</K>

290
Q

if H<0 and S>0

A

spontaneous

291
Q

if H<0 and S<0

A

spontaneous under low temp

292
Q

if H>0 and S>0

A

spontaneous under high temp

293
Q

if H>0 and S<0

A

not spontaneous

294
Q

If K>1, G is

A

negative (lnK is positive)

295
Q

gibbs sign and meaning

A

negative, spont
0, eq
positive, non-spont

296
Q

E sign and meaning

A

+, favored
-, not favored

297
Q

KE distribution graph

A

-shortest is fastest
-greater molar mass = taller + skinnier

298
Q

Don’t multiply

A

voltage

299
Q

weak means eq w/

A

water

300
Q

more magnitized

A

more unpaired electrons

301
Q

when they ask to list IMFs present, list

A

dipole-dipole

302
Q

use the word ___ when talking about X imfs

A

overcome

303
Q

IMFs are ____, not in compounds

A

between

304
Q

why more energy is required for the process occurring at vaporization than for the process occurring at fusion

A

Intermolecular attractions are completely overcome during vaporization

305
Q

if no salt bridge,

A

no voltage

306
Q

Concentration of reactant increases and concentration of product decreases, will electrochemical potential inc, dec, or remain the same?

A

Increase
Highly negative delta G indicates value of K is very large, so decreasing Q causes the system to deviate further from equilibrium, thus increasing the magnitude of the electrochemical potential

307
Q

insulation is…

A

net energy flows into wax, endothermic

308
Q

ideal vs. real gas

A

no attraction in ideal gas, real gases have attractions

309
Q

radius and nuclear charge,

A

larger radius, smaller charge

310
Q

bigger molecules have smaller

A

pressure

311
Q

when comparing moles look at

A

ratio, not amount

312
Q

qsalt =

A

-qsolution

313
Q

IMFs stronger bc

A

more polarizable electron cloud

314
Q

based on Ksp which has lowest []

A

lowest Ksp

315
Q

in a 2 step rxn easier to remove 1st H bc

A

second H is removed from a negatively charged species

316
Q

ammonium

A

NH4+

317
Q

hydroxide

A

OH-

318
Q

carbonate

A

CO3^2-

319
Q

acetate

A

CH3COO-

320
Q

chlorate

A

ClO3-

321
Q

Sulfate

A

SO4^2-

322
Q

nitrate

A

NO3-

323
Q

phosphate

A

PO4^3-

ap chem (10 decks)

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