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chem > uworld chem review > Flashcards

uworld chem review Flashcards

1
Q

covalent bond is formed by

A

valence electrons shared b/w two atoms (usually two nonmetals)

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2
Q

differences in electronegativity can result in

A

unequal electron sharing

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3
Q

in polar covalent bonds, electrons are pulled closer to the…

A

atom of greater electronegativity (giving it a partial negative charge)

atom of lesser electronegativity gets a partial positive charge

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4
Q

separation of partial charges across a polar covalent bond is…

A

a dipole moment

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5
Q

u = qr

A

q = magnitude of partial charge separated across a distance r

dipole moment equation

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6
Q

magnitude of the partial charge tends to increase as the

A

difference in electronegativity b/w the two atoms in the bond increases

bond length across which the charge is separated impacts the dipole

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7
Q

periodic trends in atomic radii allow

A

relative comparisons of bond length

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8
Q

a buffer solution is made of

A

a weak acid and its conjugate base

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9
Q

log of the ratio of base-to-acid determines

A

how much pH differs from pKa

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10
Q

a buffer system resists changes in

A

pH when OH- or H+ is added to the solution

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11
Q

buffer systems work by containing

A

large amounts of a mixture with a weak acid (HA) with its conjugate base (A-) OR

a mixture of a weak base (B) with its conjugate acid (HB+)

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12
Q

henderson-hasselbach equation is used to find

A

the pH of a buffer system

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13
Q

HA -> <- H+ + A-

A

henderson-hasselbach equation

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14
Q

acid dissociation constant

A

Ka

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15
Q

Ka =

A

[H+] [A-] / [HA]

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16
Q

pH =

A

-log [H+]

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17
Q

pKa =

A

-logKa

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18
Q

pH + henderson-hasselbach equation

A

pH = pKa + log [A-] / [HA]

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19
Q

an indicator is used to

A

visually detect the endpoint of an acid-base titation

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20
Q

an indicator is typically a

A

weak acid orbase

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21
Q

a good indicator must not react with the

A

substance being titrated and should change color close to the equivalence point of a titration

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22
Q

equivalence point is the stoichiometric amoutn of

A

titrant needed to react w/ all the solution being titrated

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23
Q

in the case of a buffer solution, a titration is…

A

halfway to teh equivalence point when a pH is achieved near the pKa value of buffer’s acid

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24
Q

bromothymol blue is a weakly acidic indiciator that changes from

A

pale yellow at pH 6 to dark blue at pH 8

only a few drops were added to the titrated solution –> little to no effect on the pH of the urine sample

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25
Q

adding a strong acid to a buffer causes the pH to

A

decrease slightly

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26
Q

adding a strong base to a buffer causes pH to

A

increase slightly

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27
Q

polyprotic acids are

A

acids w/ more than one acidic proton (H+ ion) that can dissociate in water

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28
Q

each acidic proton dissociates in a stepwise progression that generates

A

a new conjugate base with one less H+ ion until all acidic protons have dissociated

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29
Q

when a polyprotic acid is titrated with a strong base (OH-),

A

one equivalent base of base is needed to neutralize each acidic proton

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30
Q

during each step of titration, a transitory buffer system is formed as

A

some of the unreacted acid remains w/ an increasing amount of its conjugated base (formed via neutralization)

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31
Q

on a titration curve, this buffer system is indicated by a

A

plateau where the change in the pH of the solution is minimal (i.e. a buffer region)

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32
Q

titration curve of a polyprotic acid with show how many equivalent points?

A

one for each acidic proton

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33
Q

when pH = pKa

A

[HA] = [A-]

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34
Q

when pH > pKa

A

[HA] < [A-]

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35
Q

when pH < pKa

A

[HA] > [A-]

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36
Q

dissociation of an acid (HA) in water equation

A

HA + H2O -> <- H3O+ + A-

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37
Q

relative strength of an acid is described by

A

acid dissociation constant (Ka)

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38
Q

acid dissociation constant (Ka) can be also be expressed as

A

pKa

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39
Q

pKa =

A

-log Ka

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40
Q

a solution’s acidity is measured by its

A

pH

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41
Q

acidic pH depends on the

A

molar concentration of H+ ions

pH = -log[H+]

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42
Q

relationship bw pKa and pH can be described as

A

the henderson-hasselbach equation

pH = pKa + log [A-] / [HA]

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43
Q

when the pH of a solution is equal to the pKa of a weak acid…

A

equal amounts of the weak acid and its conjugate base are present in the solution

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44
Q

when pH is less than pKa

A

equilibrium is shifted towards higher concentration of the protonated acid

[HA] > [A-]

solution is more acidic

45
Q

when pH is greater than the pKa

A

equilibrium is shifted toward a higher concentration of the deprotonated conjugate base

[HA] < [A-]

46
Q

a pH of 8.0, the conjugate base HPO4^2- is the dominant because

A

pKa of H2PO4 = 6.8 (less than 8.0)

H2PO4’s pH is too high

[H2PO4-] < [HPO4 2-]

47
Q

acidic NH4+ is the dominant species in the ammonia buffer because

A

pKa of NH4+ = 9.1 (too high - must be less than 8.0 pH)

[NH4+] > [NH3]

48
Q

pKa of H2PO4 is

A

less than the pH of the solution so H2PO4- is deprotonated to form HPO4^2- as the predominant species

49
Q

pKa of NH4+ is

A

greater than solution’s pH so NH4+ remains protonated as the predominant species not NH3

50
Q

when a battery discharges, the half-rxn at the anode

A

gives up electrons (oxidation) –> transferred to the cathode (reduction to complete other half-rxn

51
Q

electron balance must be

A

present bw the half-reactions to achieve the balanced net reaction from 2 half-reactions sum

52
Q

electrons present as reactants in one half-reaction shoudl appear as

A

products in the other half-reaction

53
Q

identical species appearing as

A

reactant in one half-reaction

54
Q

products in the other half-reaction

A

cancel and are not part of the net rxn expression

55
Q

electron configuration accounts for the

A

placement of all electrons within the shells/subshells of an atom

56
Q

shells are indicated using the

A

principal quantum number (n)

relates distance of an orbital from the nucleus

57
Q

subshells are labelled by

A

type (s, p, d, f)

relates shape of occupied orbital and holds a max of 2,6, 10, 14 electrons

58
Q

electron configurations are listed

A

sequentially in order of increasing energy

59
Q

a noble gas abbreviation may be used to

A

express electron configuration more concisely

60
Q

during ionization, electrons in

A

highest energy shell/subshell are lost first

61
Q

as per the pauli exclusion principle, each orbital within

A

a subshell can hold a max of two electrons with opposite spins

62
Q

in schematic representations, electrons in subshells are represented by

A

arrows and the spin is indicated by up or down orientation

63
Q

occupied orbitals are often represented by

A

blanks or boxes

64
Q

hund’s rule states that

A

electrons fill orbitals in such a way to max number of unpaired electrons

65
Q

as a subslevel is filled, electrons form

A

pairs only after all orbitals contain at least one electron

66
Q

magnetic properties of an atom/ion depend on

A

whether or not unpaired electrons remain after all electrons are assigned to an orbita

67
Q

if unpaired electrons remain, then the atom is

A

paramagnetic and the unpaired electrons will be attracted to an external magnetic field

68
Q

if all electrons are paired then the atom is

A

diamagnetic and the paired electrons will be repelled by a magnetic field

69
Q

electron configurations of the alkaline-earth metals end with

A

a completely filled s-block

70
Q

electrons occupying the corresponding s orbital in the

A

respective valence shells are paired

71
Q

isolated forms of Ca and Sr have no unpaired electrons and are both

A

predicted to be diamagnetic

72
Q

the Ca2+ ion results from losing both electrons from the valence shell but b/c…

A

all the shells below the valence shell are fully filled (no unpaired electrons)

73
Q

in metallic solid amde of many moles of atoms…

A

the valence electrons are delocalized throughout the sample + become unpaired

= paramagnetic properties

74
Q

radioactive beta decay can occur in three forms

A

Beta-minus decay

Beta-plus decay

electron capture

75
Q

Beta-minus decay is

A

electron emission

76
Q

beta-plus decay

A

positron emission

77
Q

beta-minus decay is where a neutron converts into a

A

nuclear proton and emits an electron

78
Q

in beta-plus decay, a proton converts into a

A

neutron (opp of beta-minus decay) and emits a positron (an electron w/ a positive charge)

79
Q

in electron capture, a proton captures an electron near

A

the nucleus and converts into a neutral without a positron or electron emission

80
Q

in all three forms of beta decay, the mass number remains

A

unchanged while the atomic number increases (beta-minus decay) or decreases (beta-plus decay and electron capture) by 1

81
Q

as the atomic number changes, the identity

A

of the element changes accordingly

82
Q

if radium-226 underwent an electron capture…

A

the result would yield a nucleus of the same mass number (226) but an atomic number tha is 1 less (88 - 1 =87)

francium-226 would be detected

83
Q

when forming ionic bonds with nonmetals, the reactivity of the alkaline earth metals is observed to increasing moving down the column

A

trend in reactivity is best explained by comparing the energy required to remove an electron from each metal atom

84
Q

many atomic properties are associated with

A

electrons

85
Q

four properties associated with electrons

A

polarizability

electronegativity

electron affinity

ionization energy

86
Q

polarizability is the extent to which an

A

electron cloud can be distorted by an external charge / by an applied electric field to product a dipole

87
Q

electronegativity is the tendency of

A

an atom to attract electrons within a bond

88
Q

electron affinity assesses the tendency of an atom to

A

accept an additional electron by measuring the E change when an electron is added to an atom

89
Q

ionization energy (opp of electron affinity) measures the

A

energy required to remove an electron from an atom

90
Q

what explains how easily electrons are removed from an atom?

A

ionization energy

91
Q

moving down the alkaline-earth metal colum,

A

ionization energy decreases

makes removing an electron more favorable and increases reactivity

92
Q

although the reactivity of the alkaline-earth metals does increase as

A

electronegativity decreases

this correlation does not explain the cause of the reactivity trend

93
Q

first ionization energy refers to the

A

removal of the first electron from a neutral atom of an element (X)

94
Q

second ionization energy refers to the

A

removal of second electron

95
Q

loosely bound electrons (valence electrons) are removed

A

first before any core electrons

96
Q

removing a core electrons takes

A

more energy than does removing a valence electrons

97
Q

on the periodic table, metals generally have a lower

A

first ionization energy than nonmetals

98
Q

first ionization energies tend to increase

A

across a row, decrease down a column

99
Q

second ionization energy is the energy

A

required to remove the second of two electrons from an atom

100
Q

second ionization energy tends to increase

A

across a period and to decrease down a group

101
Q

ionization involving core electrons are

A

higher energy than those involving valence electrons

102
Q

under conditions of standard temperature and pressure (STP) are defined as

A

0 deg celsius (273 K) and 1.00 atm pressure

1.00 mole of gas occupies a volume of 22.4 L

103
Q

molar volume (only valid at STP) is often just used as

A

a known constant but this value can be derived using the ideal gas law

PV = nRT

104
Q

in a reaction, the moles of one species can be related to the

A

moles of other chemical species by applying stoichiometric mole rations obtained from a balanced rxn equation

105
Q

equilibrium is achieved in a reversible reaction when the

A

forward and reverse rxn occur simultaneously at the same rate

106
Q

at equililibrium, the forward rxn generated products as

A

fast as the reverse rxn converts those products back to OG reactants

causes concentrations of reactants and products to become constant

107
Q

although constant, the equilibrium concentrations are not necessarily

A

equal because equilibrium refers to a state of equal reaction rates (i.e. changes in concentration over time)

NOT to a state of equal concentrations

108
Q

equilibrium is achieved when

A

two opposing chem rxns occur simultaneously at the same rate so the concentrations of the chem species become constant

109
Q

equilibrium rxn rates are equal but the

A

equilibrium concentrations of chemical species may be equal

chem (2 decks)

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