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MCAT > Chem II: 1-2 > Flashcards

Chem II: 1-2 Flashcards

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MCAT flashcards
1
Q

octet rule

A

atom tends to bond with other atoms so that it has 8 electrons in outermost shell –> forming a stable electron configuration similar to that of noble gases

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2
Q

exceptions to octet rule

incomplete octet

A

stable with fewer than 8 electrons

H (2), He (2), Li (2), Be (4), B (6)

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3
Q

exceptions to octet rule

expanded octet

A

any elemet in period 3 and greater can hold more than 8 electrons

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4
Q

exceptions to octet rule

odd numbers of electrons

A

any molecule with odd number of valence electrons cannot distribute those electrons to give 8 to each atom

ex: NO

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5
Q

atoms that almost always abide by octet rule

A

C, N, O, F, Na, Mg

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6
Q

nonmetals ____ electrons

A

gain

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7
Q

metals ____ electrons

A

lose

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8
Q

ionic bonding

A

one or more electrons from an atom with low ionization energy (typically metals) are transferred to atom with high electron affinity (typically nonmetals)

resulting in electrostatic attraction between opposite chages

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9
Q

coordinate covalent

A

if both of shared electrons in covalent bond are contributed by only one of the 2 atoms

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10
Q

difference between ionic and covalent compounds

A

ionic - gain/loss of electrons

covalent - sharing of electrons

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11
Q

alkali and alkaline earth metals readily form ionic bonds with…

A

halogens of group VIIA

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12
Q

physical characteristics of ionic compounds

A
  • high MP and BP due to electrostatic attractions
  • solubility of ions in water due to interactions with polar solvents
  • good conductors of heat and electricity
  • crystal lattice arrangement to minimize repulsive forces
  • large electronegativity differences between ions
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13
Q

why do ionic bonds tend to form between metals and nonmetals

A

metals lose electrons because they have low ionization energies

nonmetals gain electrons because they have high electron affiinities

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14
Q

physical characteristics of covalent compounds

A
  • relatively week intermolecular interactions
  • lower MP and BP
  • poor conductor of electricity bc do not break down into constituent ions
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15
Q

bond length

A

avg distance between the two nuclei of atoms in a bond

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16
Q

as the number of shared electron pairs increases, bond length ______ bc…

A

decreases because the two atoms are pulled closer together

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17
Q

bond energy

A

energy required to break a bond by separating its components into their isolated, gaseous atomic states

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18
Q

the greater the number of pairs of electrons shared between the atomic nuclei, the _____ energy is required to break the bonds holding the atoms together

A

more

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19
Q

7 common diatomic molecules are:

A

(make a 7 in periodic table)

H2, N2, O2, F2, Cl2, Br2

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20
Q

dipole moment eq

A

p = qd

p: dipole moment (debye units –> coulomb meters)
q: magnitude of charge
d: distance

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21
Q

what kinds of rxns are coordinate covalent bonds usually found?

A

nucleophile-electrophile rxns, lewis acid-base rxns, complexation rxns

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22
Q

nonbonding electrons

A

electrons in valence shell that are not involved in covalent bonds

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23
Q

bonding electrons

A

valence electrons involved in a covalent bond

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24
Q

formal charge

A

assumes equal sharing of all bonded electron pairs

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25
Q

if the possible lewis structures differ in their bond connectivity or arrangement, then the lewis structures represent ______

A

different possible compounds

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26
Q

if the lewis structures show the same bond connectivity and differ only in the arrangement of the electron pairs, then these structures represent _____

A

different resonance forms of a single compounds

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27
Q

formal charge - most stable compounds

A

minimizes the number and magnitude of formal chages

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28
Q

steps in drawing lewis dot structure

A
  1. central atom - least electronegative
  2. hydrogen and halogens (F, Cl, Br, I) usually at ends
  3. usually
    1. C - 4 bonds
    2. O - double bond
    3. F - 1 bond
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29
Q

how to calculate formal charge

A

formal charge = valence electrons - nonbonding electrons - 1/2 bonding electrons

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30
Q

formal charge ____estimates the effect of electronegativity differences, while ox numbers ___estimate the effect of electronegativity differences

A

under

over

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31
Q

resonance

A

allows for greater stability, delocalizing electrons and charges over pi system

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32
Q

the more stable the structure, the ____ it contributes to the character of the resonance hybrid

A

more

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33
Q

using formal charges to assess the stability of resonance structures

lewis structures

A
  • small or no formal charges is preferred over more
  • less separation between opposite charges is preferred over larger
  • negative formal charges are placed on more electronegative atoms is more stable
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34
Q

linear

A

2 things, no lone pair

6 things, 4 lone pairs

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35
Q

trigonal planar

A

3 things, no lone pairs

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36
Q

tetrahedral

A

4 things, no lone pairs

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37
Q

trigonal bipyramid

A

5 things, no lone pairs

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38
Q

octahedral

A

6 things, no lone pair

39
Q

trigonal pyramid

A

4

1 lone pair

40
Q

seesaw

A

5

1 lone pair

41
Q

square pyramid

A

6

1 lone pair

42
Q

VSEPR?

2

A

linear

43
Q

VSEPR?

3

A

trigoonal planar

44
Q

VSEPR?

4

A

tetrahedral

45
Q

VSEPR?

5

A

trigonal bipyramid

46
Q

VSEPR?

6

A

octahedral

47
Q

VSEPR?

3

1 lone pair

A

less than 120

48
Q

VSEPR?

4

1 lone pair

A
49
Q

VSEPR?

5

1 lone pair

A

seesaw

50
Q

VSEPR?

6

1 lone pair

A
53
Q

4 things

2 lone pairs

A

<<109

bent

54
Q

5 things

2 lone pairs

A

t shaped

55
Q

5 things

3 lone pairs

A

linear

56
Q

6 things

3 lone pairs

A

t shaped

57
Q

6 things

4 lone pairs

A

linear

58
Q

6 things

2 lone pairs

A
74
Q

hybridization

3 things

A

120 deg

3sp2

1 p leftover

75
Q

hybridization

4 things

A

109 deg

4sp3

nothing left over

76
Q

hybridization

2 things

A

180 deg

2sp

2 p leftover

77
Q

electronic geometry

A

describes the spatial arrangement of all pairs of electrons around the central atom, including both the bonding and lone pairs

78
Q

molecular geometry

A

describes the spatial arrangement of only the bonding pairs of electrons

coordination number

79
Q

coordination number

A

number of atoms that surround and are bonded to a central atom

important for determining molecular geometry

80
Q

molecular orbital

A

probability of finding the bonding electron sin a given space

found by combining the wave functions of the atomic orbitals

81
Q

bonding orbital forms if

A

signs of the 2 atomic orbitals are the same

82
Q

antibonding orbital forms if

A

the signs of the two atomic orbitals are different

83
Q

sigma bond

A

when orbitals overlap head ot head

free rotation about axis

84
Q

pi bond

A

two parallel electron cloud densities

no free rotation bc cannot be twisted in a way that allows continuous overlapping of the clouds of electron densities

85
Q

In a hypothetical molecule, a Nitrogen has two sets of lone pair electrons and two covalent bonds. What would its formal charge be?

(A) -2
(B) -1
(C) 0
(D) +1

A

(B) -1

Formal Charge = 5 - [4 + 4/2] = -1

86
Q

Draw the dot structure for O3, being sure to assign formal charges as needed.

A
87
Q

Draw the dot structure for Nitrate, being sure to assign formal charges as needed.

A
88
Q

draw the three resonance structures for Nitrate.

A
89
Q

CRB In order to avoid having to draw multiple resonance structures, a resonance hybrid can be drawn. Which of the following are correct descriptions of a resonance hybrid?

I. Resonance hybrids often depict partial charges.
II. For where a double or single bond may exist, a dotted line can be used to represent the second bond.
III. Resonance hybrids are considered the “averages” of the possible resonance structures.

(A) I and II only
(B) I and III only
(C) II and III only
(D) I, II and III

A

(D) I, II and III

Each of the following are a correct description of resonance hybrids:

I. Resonance hybrids often depict partial charges.
II. For where a double or single bond may exist, a dotted line can be used to represent the second bond.
III. Resonance hybrids are considered the “averages” of the possible resonance structures.

90
Q

What is the electron geometry of Sulfur Dioxide? Molecular geometry?

A

Electron geometry of Sulfur Dioxide is trigonal planar.

Molecular geometry of Sulfur Dioxide is bent/angular.

91
Q

CRB For a tetrahedral compound, which of the following is the ideal bond angle (in degrees)?

(A) 90
(B) 104.5
(C) 109.5
(D) 120

A

(C) 109.5

For a tetrahedral compound, the ideal bond angle is 109.5 degrees.

92
Q

What is the electron geometry of NH3? Molecular geometry?

A

Electron geometry of NH3 is tetrahedral.

Molecular geometry of NH3 is trigonal pyramidal.

93
Q

What is the electron geometry of H2O? Molecular geometry?

A

Electron geometry of H2O is tetrahedral.

Molecular geometry of H2O is bent.

94
Q

What is the Electron Geometry of ICl2+? Molecular Geometry?

A

Electron Geometry of ICl2+ is Tetrahedral.

Molecular Geometry of ICl2+ is Bent.

95
Q

a) aluminum
b) calcium
c) vanadium
d) scandium

A

c) vanadium

96
Q

a) being in very polar solvents
b) having component atoms that have an odd number of valence electrons in their elemental forms
c) being in polar aprotic solvents
d) having an odd number of valence electrons across the entire molecule

A

d) having an odd number of valence electrons across the entire molecule

97
Q
A

c

98
Q
A

b) PCl3

99
Q
A

a

100
Q
A

a

101
Q
A

c) 2

102
Q
A

b) 0

103
Q
A

c) 16

104
Q

intermolecular forces from weakest to strongest

A

london disperson forces/van der waals < dipole-dipole interactions < hydrogen bonding

105
Q

london dispersion forces

A
  • type of van der waals force
  • rapid polarization and counterpolarization of electron cloud and formation of short-lived dipole moments
  • weakest of intermolecular interactions
  • do not extend over long distances
  • relevant only when molecules are in close proximity
  • strength depends on degree and ease by which molecules can be polarized
    • large molecules more easily polarizable
106
Q

dipole dipole interactions

A
  • present in solid and liquid phases
    • negligible in gas phase bc of inc distance between gas particles
  • why polar species have higher MP and BP
107
Q

hydrogen bonding

A
  • strongest
  • when H is bonded to NOF (very electronegative)
  • may be intra or inter moelcualr
  • no actual sharing or transferring of electrons between two atoms
  • unusually high BP and MP
108
Q

What is the Electron Geometry of SF4? Molecular Geometry?

A

Electron Geometry of SF4 is Trigonal Bipyramidal.

Molecular Geometry of SF4 is See-saw.

109
Q

What are the bond angles for Trigonal Bipyramidal? See-saw?
T-structure? Linear?

A
110
Q

What are the bond angles for Octahedral? Square Pyramidal? Square Planar?

A
111
Q

a) ClF3
b) H2O
c) CH4
d) XeF2

A

d) XeF2

MCAT (26 decks)

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