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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
ionization energy
the minimum energy needed to remove an electron from an atom or ion in the gas phase
26
second ionization energy will be (smaller/larger) than first
larger (think Zeff)
27
electron affinity
the energy associated with the addition of an electron to the valence shell of an atom that is in the gas phase
28
electron affinity periodic trend
becomes more negative across a period (increases) and no definite trend for down a group
29
metal characteristics
malleable, ductile, shiny, conduct heat/electricity, form cations, oxidized
30
nonmetal characteristics
brittle, dull, electrical/thermal insulators, form anions, reduced
31
metallic character
how closely an element's properties match the ideal properties of a metal
32
metallic character periodic trend
decreases across a period and increases down a group
33
ionic bonding
metal & nonmetal
34
covalent bonding
nonmetal (becomes the cation) & nonmetal (becomes the anion)
35
ionic bonds involve the
transfer of electrons from one atom to another
36
covalent bonds involve the
sharing of electrons between two atoms
37
molecular compounds are composed of
atoms covalently bonded to each other
38
ionic compounds are composed of
ionic bonds
39
chemical formula
represents a compound; indicates the type and number of each electron present in the compound
40
chemical formulas can generally be categorized into three different types
empirical, molecular, structural
41
empirical formula
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
42
molecular formula
gives the actual number of atoms of each element in a molecule of a compound
43
structural formula
a sketch or diagram of how the atoms in the molecule are bonded to each other
44
octet rule
when atoms bond, they tend to gain, lose, or share e-s to result in a noble gas-like configuration
45
expanded octets
they involve the nonmetal elements located in period 3 and below
46
ionic compounds can be categorized into two types
metal forms only one type of ion & metal forms more than one type of ion
47
naming binary ionic compounds of type I cations
(name of cation (metal)) + (base name of anion (nonmetal) + -ide)
48
naming type II binary ionic compounds
(name of the cation (metal)) + (charge of the cation in roman numerals in parentheses) + (base name of anion + -ide)
49
oxyanions
anions containing oxygen and another element
50
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 ?
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
hydrates
ionic compounds containing a specific number of water molecules associated with each formula unit
52
hydrate prefixes
``` hemi 1/2 mono 1 di 2 tri 3 tetra 4 penta 5 hexa 6 hepta 7 octa 8 ```
53
bonding pairs
electrons that are shared by atoms
54
lone/nonbonding pairs
electrons that are not shared by atoms but belong to a particular atom
55
naming binary molecular compounds
prefix + name of 1st element + prefix + (base name of 2nd element + -ide)
56
prefixes for molecular compounds
``` mono 1 di 2 tri 3 tetra 4 penta 5 hexa 6 hepta 7 octa 8 nona 9 deca 10 ```
57
molecular mass
the mass of an individual molecule or formula unit; the mass of one mole of that compound
58
formula mass
(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
percent mass of an element
(molecular mass of element Z) / (mass of 1 mole of compound) x 100%
60
finding an empirical formula
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
from empirical to molecular formula
(empirical formula) * n, where n = molar mass / empirical formula molar mass
62
combustion analysis
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
metallic bond
metal and metal
64
electronegativity
the relative ability of an atom to attract electrons in a bond to itself
65
electronegativity periodic trend
increases across a period and decreases down a group
66
the larger the difference in ?, the more polar the bond
electronegativity
67
ΔEN 0-0.4
``` 0 = pure covalent 0.1-0.4 = nonpolar covalent ```
68
ΔEN 0.4-2.0
polar covalent
69
ΔEN 2.0+
ionic
70
dipole moment (μ)
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
percent ionic character
the % of a bond's measured dipole moment compared to what it would be if the electrons were completely transferred; increases with ΔEN
72
resonance
used when two or more valid Lewis structures can be drawn for the same compound
73
formal charge
an electron bookkeeping system that allows us to discriminate between alternative Lewis structures
74
rules of resonance structures
same connectivity, same number of electrons, second row elements have a max of 8 e-s, formal charges must total the same
75
formal charge rules
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
central atom in Lewis structures should me (least/most) EN
least
77
bond energy
the amount of energy, in the gaseous state, that it takes to break one mole of a bond in a compound
78
trends in bond energies
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
bond length
the distance b/w the nuclei of bonded atoms
80
trends in bond length
decreases across a period and increases down the column
81
molecular geometry
describes the shape of a molecule with terms that relate to geometric figures
82
VSEPR (Valence Shell Electron Pair Repulsion) Theory
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
electron groups
each long pair of e-s, each bond (regardless of whether it is single, double, or triple)
84
two e- groups electron geometry
linear, 180*
85
three e- groups electron geometry
trigonal planar, 120*
86
four e- groups electron geometry
tetrahedral, 109.5*
87
five e- groups electron geometry
trigonal bipyramidal, 90* axial 120* equatorial
88
six e- groups electron geometry
octahedral, 90*
89
relative sizes of repulsive force interactions
bonding pair to bonding pair < lone pair to bonding pare < lone pair to lone pair
90
lone pairs affect the ?
molecular geometry
91
using VSEPR to predict molecular geometries
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
for a molecule to be polar, it must have
polar bonds and an unsymmetrical shape
93
2 e- groups, no lone pairs molecular geometry
linear
94
3 e- groups, no lone pairs molecular geometry
trigonal planar
95
3 e- groups, 1 lone pair molecular geometry
bent
96
4 e- groups, no lone pairs molecular geometry
tetrahedral
97
4 e- groups, 1 lone pair molecular geometry
trigonal pyramidal
98
4 e- groups, 2 lone pairs molecular geometry
bent
99
5 e- groups, no lone pairs molecular geometry
trigonal bipyramidal
100
5 e- groups, 1 lone pair molecular geometry
seesaw
101
5 e- groups, 2 lone pairs molecular geometry
T-shaped
102
5 e- groups, 3 lone pairs molecular geometry
linear
103
6 e- groups, no lone pairs molecular geometry
octahedral
104
6 e- groups, 1 lone pair molecular geometry
square pyramidal
105
6 e- groups, 2 lone pairs molecular geometry
square planar

Fall 2017 (24 decks)

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