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Fall 2017 > Chem 105 Test 2 (Ch. 3-5) > Flashcards

Chem 105 Test 2 (Ch. 3-5) Flashcards

1
Q

main-group elements

A

groups 1-2 (s orbitals)

groups 13-18 (p orbitals)

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2
Q

transition elements

A

groups 3-12 (d orbitals)

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3
Q

inner transition elements

A

f orbitas

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4
Q

the Pauli Exclusion principle

A

no two electrons in an atom may have the same set of four quantum numbers

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5
Q

the sublevels (s, p, d, f, etc.) in each principal energy shell of hydrogen, or other single electron systems, all have (the same/different) energy, so they are said to be ?

A

the same; degenerate

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6
Q

for multi-electron atoms, the energies of the sublevels are split due to

A

charge interaction, shielding, and penetration

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7
Q

aufbau principle

A

electrons fill atomic orbitals from lowest energy to highest

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8
Q

Coulomb’s Law

A

for like charges, the PE is positive and decreases as the particles get farther apart (r increases)
for opposite charges, the PE is negative and becomes more negative as the particles get closer together
the strength of the interaction increases as the size of the charges increases

essentially, E is proportional to q1q2 / r

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9
Q

shielding

A

when repulsion from other electrons in the nucleus cause an electron to experience a net reduction in attraction to the nucleus

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10
Q

effective nuclear charge (Zeff)

A

the total amount of attraction that an electron feels for the nucleus’s protons

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11
Q

attraction between nucleus protons and orbiting electrons is related to

A

the orbital type the electron occupies (ex: e-s in the s orbital are better shielders than e-s in p orbitals - s>p>d>f)

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12
Q

the degree of penetration is related to

A

the orbital’s radial distribution function

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13
Q

penetration causes the energies of sublevels in the same principal level to

A

not be degenerate

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14
Q

Hund’s rule

A

when filling orbitals that have the same energy (degenerate), place one electron in each orbital before completing pairs

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15
Q

electrons in lower-energy shells

A

core (inner) electrons

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16
Q

electrons in all the sublevels with the highest principal energy level are called the

A

valence electrons

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17
Q

one of the most important factors in the way an atom behaves, both chemically and physically, is ? because

A

the number of valence e-s because the valence e-s participate in bonding & the valence shell is where e-s are lost or added to make cations/anions

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18
Q

some transition metals have irregular electron configurations in which the ns only partially fills before the (n-1)d or completely filled sublevel - which do we need to know?

A

Cr, Mo (half-filled sublevel)

Cu, Ag (completely filled sublevel)

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19
Q

Zeff periodic trend

A

increases as you go across a period and decreases as you go down a column

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20
Q

Zeff is calculated by

A

Z - S

Z = nuclear charge, S = number of electrons in lower energy levels

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21
Q

atomic radii periodic trend

A

decreases across a period and increases down a group

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22
Q

atomic radii for transition elements

A

increases down the column but roughly the same across the d block

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23
Q

cation radius is (smaller/larger) than atom radius and anion radius is (smaller/larger) than atom radius

A

smaller, larger

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24
Q

ionization energy periodic trend

A

decreases down a group and increases across the period

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25
Q

ionization energy

A

the minimum energy needed to remove an electron from an atom or ion in the gas phase

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26
Q

second ionization energy will be (smaller/larger) than first

A

larger (think Zeff)

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27
Q

electron affinity

A

the energy associated with the addition of an electron to the valence shell of an atom that is in the gas phase

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28
Q

electron affinity periodic trend

A

becomes more negative across a period (increases) and no definite trend for down a group

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29
Q

metal characteristics

A

malleable, ductile, shiny, conduct heat/electricity, form cations, oxidized

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30
Q

nonmetal characteristics

A

brittle, dull, electrical/thermal insulators, form anions, reduced

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31
Q

metallic character

A

how closely an element’s properties match the ideal properties of a metal

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32
Q

metallic character periodic trend

A

decreases across a period and increases down a group

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33
Q

ionic bonding

A

metal & nonmetal

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34
Q

covalent bonding

A

nonmetal (becomes the cation) & nonmetal (becomes the anion)

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35
Q

ionic bonds involve the

A

transfer of electrons from one atom to another

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36
Q

covalent bonds involve the

A

sharing of electrons between two atoms

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37
Q

molecular compounds are composed of

A

atoms covalently bonded to each other

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38
Q

ionic compounds are composed of

A

ionic bonds

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39
Q

chemical formula

A

represents a compound; indicates the type and number of each electron present in the compound

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40
Q

chemical formulas can generally be categorized into three different types

A

empirical, molecular, structural

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41
Q

empirical formula

A

gives the relative number of atoms of each element in a compound; the simplest whole number ratio representation of the group and number of elements present in a molecule

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42
Q

molecular formula

A

gives the actual number of atoms of each element in a molecule of a compound

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43
Q

structural formula

A

a sketch or diagram of how the atoms in the molecule are bonded to each other

44
Q

octet rule

A

when atoms bond, they tend to gain, lose, or share e-s to result in a noble gas-like configuration

45
Q

expanded octets

A

they involve the nonmetal elements located in period 3 and below

46
Q

ionic compounds can be categorized into two types

A

metal forms only one type of ion & metal forms more than one type of ion

47
Q

naming binary ionic compounds of type I cations

A

(name of cation (metal)) + (base name of anion (nonmetal) + -ide)

48
Q

naming type II binary ionic compounds

A

(name of the cation (metal)) + (charge of the cation in roman numerals in parentheses) + (base name of anion + -ide)

49
Q

oxyanions

A

anions containing oxygen and another element

50
Q

oxyanions are named according to the number of oxygen atoms in the ion: if there are two ions in the series ? and if there are more than two ions in the series ?

A

if two, the one with less oxygen atoms has the ending -ite and the one with more has the ending -ate
if more than two, hypo- -ite, -ite, -ate, per- -ate

51
Q

hydrates

A

ionic compounds containing a specific number of water molecules associated with each formula unit

52
Q

hydrate prefixes

A 
hemi 1/2
mono 1
di 2
tri 3
tetra 4
penta 5
hexa 6
hepta 7
octa 8
53
Q

bonding pairs

A

electrons that are shared by atoms

54
Q

lone/nonbonding pairs

A

electrons that are not shared by atoms but belong to a particular atom

55
Q

naming binary molecular compounds

A

prefix + name of 1st element + prefix + (base name of 2nd element + -ide)

56
Q

prefixes for molecular compounds

A 
mono 1
di 2
tri 3
tetra 4
penta 5
hexa 6
hepta 7
octa 8
nona 9
deca 10
57
Q

molecular mass

A

the mass of an individual molecule or formula unit; the mass of one mole of that compound

58
Q

formula mass

A

(number of atoms of 1st element in chemical formula x atomic mass of 1st element) + (number of atoms of 2nd element in chemical formula x atomic mass of 2nd element) + …

59
Q

percent mass of an element

A

(molecular mass of element Z) / (mass of 1 mole of compound) x 100%

60
Q

finding an empirical formula

A
  1. convert the percentage to grams
  2. convert grams to moles
  3. divide all by the smallest number of moles to obtain the atom-to-atom ratio for each of the elements in the compound
  4. multiply all mole ratios by a number to make all whole numbers
61
Q

from empirical to molecular formula

A

(empirical formula) * n, where n = molar mass / empirical formula molar mass

62
Q

combustion analysis

A

a common technique for analyzing compounds is to burn a known mass of compound and weigh the amounts of product made; by knowing the mass of the product and composition of constituent element in the product, the original amount of constituent element can be determined

63
Q

metallic bond

A

metal and metal

64
Q

electronegativity

A

the relative ability of an atom to attract electrons in a bond to itself

65
Q

electronegativity periodic trend

A

increases across a period and decreases down a group

66
Q

the larger the difference in ?, the more polar the bond

A

electronegativity

67
Q

ΔEN 0-0.4

A 
0 = pure covalent
0.1-0.4 = nonpolar covalent
68
Q

ΔEN 0.4-2.0

A

polar covalent

69
Q

ΔEN 2.0+

A

ionic

70
Q

dipole moment (μ)

A

a measure of bond polarity; directly proportional to the size of the partial charges (q) and the distance (r) between them
μ = q*r
measured in Debyes, D

71
Q

percent ionic character

A

the % of a bond’s measured dipole moment compared to what it would be if the electrons were completely transferred; increases with ΔEN

72
Q

resonance

A

used when two or more valid Lewis structures can be drawn for the same compound

73
Q

formal charge

A

an electron bookkeeping system that allows us to discriminate between alternative Lewis structures

74
Q

rules of resonance structures

A

same connectivity, same number of electrons, second row elements have a max of 8 e-s, formal charges must total the same

75
Q

formal charge rules

A
  1. sum in neutral atom must be zero
  2. sum in an ion must equal the charge of the ion
  3. small (or zero) formal charges on individual atoms are better than large ones
  4. when formal charge cannot be avoided, negative FC should reside on the most EN negative
76
Q

central atom in Lewis structures should me (least/most) EN

A

least

77
Q

bond energy

A

the amount of energy, in the gaseous state, that it takes to break one mole of a bond in a compound

78
Q

trends in bond energies

A

in general, the more e-s two atoms share, the stronger the covalent bond
in general, the shorter the covalent bond, the stronger the bond (atomic radii)

79
Q

bond length

A

the distance b/w the nuclei of bonded atoms

80
Q

trends in bond length

A

decreases across a period and increases down the column

81
Q

molecular geometry

A

describes the shape of a molecule with terms that relate to geometric figures

82
Q

VSEPR (Valence Shell Electron Pair Repulsion) Theory

A

e- groups around the central atoms will be most stable when they are as far apart as possible; the resulting geometric arrangement will allow us to predict the shapes and bond angles in the molecule

83
Q

electron groups

A

each long pair of e-s, each bond (regardless of whether it is single, double, or triple)

84
Q

two e- groups electron geometry

A

linear, 180*

85
Q

three e- groups electron geometry

A

trigonal planar, 120*

86
Q

four e- groups electron geometry

A

tetrahedral, 109.5*

87
Q

five e- groups electron geometry

A

trigonal bipyramidal, 90* axial 120* equatorial

88
Q

six e- groups electron geometry

A

octahedral, 90*

89
Q

relative sizes of repulsive force interactions

A

bonding pair to bonding pair < lone pair to bonding pare < lone pair to lone pair

90
Q

lone pairs affect the ?

A

molecular geometry

91
Q

using VSEPR to predict molecular geometries

A
  1. draw the Lewis structure
  2. determine the # e- groups around the central atom
  3. classify each e- group as bonding or lone pair and count each type
  4. use table 5.5 to det. shape & bond angles
92
Q

for a molecule to be polar, it must have

A

polar bonds and an unsymmetrical shape

93
Q

2 e- groups, no lone pairs molecular geometry

A

linear

94
Q

3 e- groups, no lone pairs molecular geometry

A

trigonal planar

95
Q

3 e- groups, 1 lone pair molecular geometry

A

bent

96
Q

4 e- groups, no lone pairs molecular geometry

A

tetrahedral

97
Q

4 e- groups, 1 lone pair molecular geometry

A

trigonal pyramidal

98
Q

4 e- groups, 2 lone pairs molecular geometry

A

bent

99
Q

5 e- groups, no lone pairs molecular geometry

A

trigonal bipyramidal

100
Q

5 e- groups, 1 lone pair molecular geometry

A

seesaw

101
Q

5 e- groups, 2 lone pairs molecular geometry

A

T-shaped

102
Q

5 e- groups, 3 lone pairs molecular geometry

A

linear

103
Q

6 e- groups, no lone pairs molecular geometry

A

octahedral

104
Q

6 e- groups, 1 lone pair molecular geometry

A

square pyramidal

105
Q

6 e- groups, 2 lone pairs molecular geometry

A

square planar

Fall 2017 (24 decks)

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