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From Flashcardlet > Chem Test 2 > Flashcards

Chem Test 2 Flashcards

1
Q

oxidation number

A

the charge an atom would have if e- were completely transferred (like in an ionic compound)- book keeping method to track electrons

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2
Q

Rules for oxidation numbers

A
  • in elemental form (Cl2, O2, Na)- O.N=0
  • for monoatomic ions- O.N= ion charge
  • sum of O.N for atoms in a compound- O.N=0
  • sum of O.N for atoms in a polyatomic ion- O.N=ion charge
  • Group 1A Alkali Metal- O.N=+1
  • Group 2A Alkaline Earth Metals- O.N=+2
  • Hydrogen-O.N= +1 (w/ nonmetals)
  • Fluorine- O.N= -1
  • Oxygen- O.N= -2
  • Group 7A Halogens-O.N= -1 (except w/ more EN)
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3
Q

bond length

A

some distance attracts balance repulsion, is about the sum of atomic radii

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4
Q

bond energy and length

A

max attraction and min repulsion is the best distance

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5
Q

shortest bond= ___

A

triple, most energy to break

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6
Q

longest bond= ____

A

single, least energy to break

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7
Q

polar covalent bond

A

not equal sharing of e-, greater e- density around one atom

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8
Q

E.N.D

A

electronegativity difference, ability to hold e-

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9
Q

END= 0

A

nonpolar, pure covalent

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10
Q

END 0<2

A

polar covalent

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11
Q

END > 2

A

ionic

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12
Q

formal charge

A

difference btwn # of valence e- in an isolated atom and the # of e- assigned to that atom in a Lewis Structure; “charge an atom would have if all e- were equally shared”

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13
Q

Formal charge formula

A

FC=VE – NBE – 1/2BE

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14
Q

resonance structure

A

one of 2+ lewis structures for a single molecule that cannot be represented w/ 1 structure

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15
Q

bond order

A

of bonds

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16
Q

delocalized

A

electrons can move, they are spread out over space, represented by —–

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17
Q

incomplete octet

A

exception to octet rule, need valence e- to form octet, ex. BeH2

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18
Q

odd e- molecules

A

exception to octet rule, use formal charges to dictate which is best, ex. NO

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19
Q

expanded octet

A

exception to octet rule, principal quantum # n>2 (n=3+), ex. SF6

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20
Q

VSEPR

A

valence shell electron pair repulsion. Minimize repulsions each group of electrons are located as far as possible from the other e-

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21
Q

2 groups

A

180, linear

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22
Q

3 groups

A

120, trigonal planar

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23
Q

4 groups

A

109.5, tetrahedral

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24
Q

5 groups

A

90 and 120, trigonal bipyramidal

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25
Q

6 groups

A

90, octahedral

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26
Q

factors that affect bond angle:

A

2 lone pairs>lone pair & bonding pair> 2 bonding pair

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27
Q

bond angle

A

angle formed by the nuclei of 2 surrounding atoms with the central atom nucleus

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28
Q

valence bond theory

A
  • a covalent bond forms when the orbitals of 2 atoms overlap and a pair of electrons occupy the overlap region
  • depends on orbital shape and direction
  • the greater the overlap, the stronger the bond
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29
Q

hybridization

A

the concept of mixing atomic orbitals into new hybrid orbitals (with different energies, shapes, etc., than the component atomic orbitals) suitable for the pairing of electrons to form chemical bonds in valence bond theory
way of correcting bond angles

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30
Q

linear hybrid

A

sp

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31
Q

trigonal planar hybrid

A

sp2

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32
Q

tetrahedral hybrid

A

sp3

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33
Q

trigonal bipyramidal

A

sp3d

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34
Q

octahedral

A

sp3d2

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35
Q

sigma

A

orbital overlap end to end; ALL SINGLE BONDS

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36
Q

pi

A

sideways overlap, weaker; DOUBLE AND TRIPLE BONDS

37
Q

when do double bonds occur?

A

when e- are placed in unhybridized orbitals, contain sigma and pi bond

38
Q

what can/cannot rotate?

A

can- single

cannot- double, triple

39
Q

cis isomer

A

same side of double bond

40
Q

trans isomer

A

opposite side of double bond

41
Q

aufbau principle

A

e- will fill lowest orbital first

42
Q

hund’s rule

A

most stable is not paired, instead parallel spin

43
Q

pauli exclusion principle

A

no 2e- will have the same 4 quantum numbers

44
Q

bond order

A

1/2(number of bonding e-) – (number of anti-bonding e-)

45
Q

homo

A

highest occupied molecular orbital

46
Q

lumo

A

lowest unoccupied molecular orbital

47
Q

bond polarity

A

greater e- density around one atom, if elements are different= polar, if same=nonpolar

48
Q

MO diagram rules

A

e- MO=#e- AO
Stable bonding=less stable antibonding
Fill orbitals Low E -> High E
Follow aufbau, hund, pauli

49
Q

molecular polarity depends on

A

shape and bond polarity

50
Q

a molecule is polar if… (2 things)

A

has one or more polar bonds

individual bond dipoles do not cancel

51
Q

dipole moment

A

mathematical measurement of polarity

52
Q

in an electric field polar molecules..

A

align themselves with the electricity of the field

53
Q

gas

A

particles are spaced far apart, freely move, no fixed volume, take shape of their container

54
Q

liquid

A

closer together than a gas, can move around, take shape of their container

55
Q

solid

A

very close together, no change of shape, can vibrate but not move

56
Q

2 types of molecular forces

A

intramolecular and intermolecular

57
Q

intramolecular forces

A

WITHIN bonds; the physical bond itself

58
Q

intermolecular

A

BTWN molecules; melting point, boiling point etc

59
Q

internuclear axis

A

sigma bonds always have this, bond btwn nuclei of atoms

60
Q

LCAO

A

linear combination of atomic orbitals

61
Q

what does MO theory state?

A

orbitals are a combination of atomic orbital wave functions

62
Q

London Dispersion Forces are in _____?

A

EVERYTHING; polar and nonpolar

63
Q

what IMFs are nonpolar molecules made of?

A

London Dispersion Forces

64
Q

what is the weakest IMF

A

LDF, caused by fluctuations in e- distribution

65
Q

instantaneous dipole

A

e- move to one side and create partial positive and negative charges

66
Q

polarizability

A

the ease with which electron distribution in the atom/molecules can be distorted

67
Q

what increases polarizability?

A

greater # of electrons (higher molar mass)

more diffuse an electron cloud (less electronegative)

68
Q

dispersion forces increases as ____ increases

A

molar mass

69
Q

dipole-dipole forces

A

between polar molecules, stronger than LDF

70
Q

Hydrogen bonds

A

type of dipole-dipole interaction btwn a H and an EN O,N or F which can interact with a different O,N or F in another molecule; NOT A REAL BOND, just VERY STRONG

71
Q

strongest IMF

A

Hydrogen

72
Q

dipole dipole is btwn what

A

polar molecules

73
Q

evidence for H bonds?

A

H2O, HF, and NH3 are light yet have high boiling point because they contain H bonds, which are strong and harder to break

74
Q

which form of water is denser? Solid or liquid?

A

Liquid, @ 4 degrees Celsius

75
Q

STP

A

standard temperature and pressure
T= 0C or 273K
P= 1 atm or 760 Torr

76
Q

Standard Molar Volume of a Gas

A

22.4 L

77
Q

Charles Law

A

V1/T1 = V2/T2

78
Q

Boyles Law

A

P1V1=P2V2

79
Q

Avogadros Law

A

V1/n1=V2/n2

80
Q

Combined Gas Law

A

P1V1/T1=P2V2/T2

81
Q

Ideal Gas Law

A 
PV=nRT
P in atm
V in liters
n is moles
R=0.0821
T in Kelvin
82
Q

how to find concentration

A

P=[A]RT or [A]=n/V

83
Q

mole fraction

A

Xi=ni/nT

84
Q

partial pressure

A

Pi=XiPT
Ptotal=P1 + P2
PA= nART/V

85
Q

kinetic energy equation

A

KE=1/2mv2

about=RT when R=8.314 J/mol*K

86
Q

at the same temperature, ligth gases will have a ___ ___ than heavier gases

A

faster speed

87
Q

at the same temperature kinetic energies are ___

A

equal

88
Q

what are gases particle collisions?

A

perfectly elastic

89
Q

gases do not ____ to each other

A

exert attractive or repulsive forces

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