Exam 2 Flashcards

1
Q

principle quantum number

A

n = energy level of electron/orbital
positive integer

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2
Q

angular momentum quantum number

A

l = shape of orbital
l = 0, 1, 2, …, n-1

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3
Q

magnetic quantum number

A

mₗ = orbital orientation
mₗ = -l, …, 0, …, l

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4
Q

spin quantum number

A

mₛ = spin of specific electron
mₛ = ±1/2

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5
Q

letter names associated with l values
l = 0, 1, 2, 3, 4

A

l = s, p, d, f, g

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6
Q

sketch orbital:
n=1
l=0

A

1s
0 nodes - nucleus doesn’t count
sphere

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7
Q

sketch orbital:
n=2
l=0

A

2s
1 node - radial
hollow sphere

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8
Q

sketch orbital:
n=2
l=1

A

2p
1 node - planar
dumbbell shape

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9
Q

sketch orbital:
n=3
l=1

A

3p
2 nodes - 1 planar, 1 radial
dumbbell with 2 hats

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10
Q

sketch orbital:
n=3
l=2

A

3d
2 nodes - planar
4 leaf clover or dumbbell with donut

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11
Q

sketch orbital:
n=4
l=2

A

4d
3 nodes - 2 planar, 1 radial
4 leaf clover or dumbbell with 4 hats

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12
Q

Hund’s Rule

A

Electrons should be placed evenly across degenerate orbitals before doubling up.

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13
Q

Pauli’s Exclusion Principle

A

Every electron within an atom has a unique identity, as determined by its n, l, mₗ, and mₛ values.

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14
Q

Aufbau Principle

A

When determining which electrons are in which orbitals, start from the lowest energy possibilities (orbitals).

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15
Q

paramagnetic

A

there are unpaired electrons in the atom
net spin non-zero
affected by magnetic fields

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16
Q

diamagnetic

A

all electrons are paired within the atom
net spin = 0

17
Q

atomic radius
(definition and periodic trends)

A

half of the distance between 2 bonded of an element
decreases across periods
increases down groups

18
Q

effective nuclear charge
(definition and periodic trends)

A

strength of the pull of the nucleus on electrons
increases across periods
decreases down groups

19
Q

ionization energy
(definition, periodic trends, and first vs. second ionization energy)

A

the energy required to move a valence electron
increases across periods
decreases down groups
first vs. second….

20
Q

ionic bonds nomenclature
e.g. NaCl
e.g. MgSO₄
e.g. FeCl₃

A
  • metal (cation), then nonmetal (anion)
  • cation element name + anion with “-ide” ending
  • add roman numeral after cation if needed
    e.g. sodium chloride
    e.g. magnesium sulfate
    e.g. iron (III) chloride
21
Q

cation nomenclature
e.g. Al⁺
e.g. Fe³⁺

A

add “ion”
add roman numeral after element name if transition metal
e.g. aluminum ion
e.g. iron (III) ion

22
Q

monoatomic anion nomenclature
e.g. Cl⁻
e.g. O²⁻

A

add suffix -ide
e.g. chloride
e.g. oxide

23
Q

molecular compound nomenclature PREFIXES
1-10

A

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

24
Q

molecular compound nomenclature
e.g N₂O
e.g. ICl₃

A
  • less electronegative element first
  • add prefixes except for mono- on first element
  • add “-ide” suffix to last element
    e.g. dinitrogen monoxide
    e.g. iodine trichloride
25
electronegativity (definition, periodic trends, polarity of molecular compounds)
the ability of an atom to attract bonded electrons to itself increases across periods decreases down groups polar bonds occur when there is an electronegativity difference -- polar compounds exist when vectors don't cancel
26
Formal charge (in a Lewis structure)
Calculate formal charge by subtracting the number of valence electrons around it in a compound from its atomic number - includes all lone pair electrons and half of all bonded electrons Prioritize lowest formal charges on each atom and opposite charges next to each other
27
Violations of octet rule (Lewis structures)
- odd number of valence electrons --> free radical - incomplete octet --> usually caused by formal charge - expanded octet (period 3 or below) --> more electrons on central atom
28
VSEPR electron domain geometries and angles for 2-6 domains
2 domains = linear (180°) 3 domains = trigonal planar (120°) 4 domains = tetrahedral (109.5°) 5 domains = trigonal bipyramidal (120° and 90°) 6 domains = octahedral (90°)
29
2 electron domains: molecular geometries
2 bonding = linear
30
3 electron domains: molecular geometries
2 bonding = bent (<120°) 3 bonding = trigonal planar (120°)
31
4 electron domains: molecular geometries
2 bonding = bent (<<109.5°) 3 bonding = trigonal pyramidal (<109.5°) 4 bonding = tetrahedral (109.5°)
32
5 electron domains: molecular geometries + bond positions
2 bonding (axial positions) = linear (180°) 3 bonding (axial + one equatorial) = T-shaped (90°) 4 bonding (axial + two equatorial) = seesaw (90° and <120°) 5 bonding = trigonal bipyramidal (90° and 120°)
33
6 electron domains: molecular geometries (start at 4 bonding)
4 bonding = square planar (90°) 5 bonding = square pyramidal (<90°) 6 bonding = octahedral (90°)
34
valence bond theory (2 types of bonds)
sigma (σ) bonds: head-to-head overlap --> single bonds pi (π) bonds: side-to-side overlap --> double bonds
35
valence bond theory (significance of single, double, and triple bonds)
single bond = 1 σ bond double bond = 1 σ bond and 1 π bond triple bond = 1 σ bond and 2 π bonds
36