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unit 3 Flashcards

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1
Q

When two atoms approach each other repulsions between each atom’s:

A

negatively charged electron clouds

Positively charged nuclei

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2
Q

Attractions between nucleus and approaching atoms electron cloud are strongest where

A

electron clouds overlap between adjacent nuclei

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3
Q

If attractive forces stronger than repulsive forces

A

Atoms have a lower energy than when apart

Chemical bond forms

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4
Q

ionic bond

A

electrostatic attraction between a metal cation and nonmetal anion

EN difference over 1.7

transfer of electrons

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5
Q

crystal lattice

A

oppositely charged ions form an ordered, solid, 3-D array with large numbers of interionic forces

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6
Q

crystal lattice properties

A

High melting and boiling points

Chemical formula of an ionic compound is the smallest whole-number ratio of ions

Different crystal structures depend on sizes and ratios of ions

Ratios depend on ionic charges in the compound

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7
Q

describe what happens when 2 H atoms approach each other

A

Electron cloud of one atom is attracted to nucleus of other atom, kinetic energy increases

Repulsive forces as nuclei approach each other slows atoms, kinetic energy becomes potential energy

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8
Q

what happens when electron clouds overlap?

A

Attractive forces exceed repulsive forces

Valence electrons move into space between the two nuclei where there is most attractive force between nuclei

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9
Q

covalent bond

A

sharing of electrons between 2 nonmetals

attraction between a pair of electrons and two nuclei

usually independent molecules

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10
Q

non polar covalent bond

A

EN difference = 0-0.3

atoms of the same element bonded

electrons equidistant between nuclei

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11
Q

polar covalent bond

A

EN difference = 0.4-1.6

bonding electrons pulled closer to more electronegative atom

partially positive end has less electron density

partially negative end has more electron density

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12
Q

number of electrons needed =

A

number of bonds formed in a covalent compound

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13
Q

ionic compound properties

A

Strong electrostatic attractions between charged ions must be broken to melt-high melting point

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14
Q

covalent compounds properties

A

Don’t always need to break bonds between atoms

Weak intermolecular attractions need to be broken to melt

Separation of molecules not breaking of bonds between atoms

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15
Q

network covalent solids

A

solid where all the atoms are covalently bonded in a continuous network

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16
Q

network covalent solids properties

A

Covalent bonds extend through entire sample

High melting points

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17
Q

For a chemical bond to form between two atoms:
How must the energy associated with the bonded atoms compare to the energy when the atoms are apart?

A

The energy of the bonded atoms must be lower than the energy of the atoms when they are apart

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18
Q

For a chemical bond to form between two atoms: What does this tell us about the attractive forces compared to the repulsive forces between them?

A

The attractive forces between the bonded atoms are stronger than the repulsive forces

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19
Q

What is an ionic crystal lattice and how does it explain the high melting points of ionic compounds?

A

An ionic crystal lattice is an ordered, solid, three-dimensional array of cations and anions

The large number of interionic forces in the crystal lattice locks the ions in place giving ionic compounds their high melting points

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20
Q

What are the attractive forces associated with
Ionic bonds

A

Electrostatic attractions between oppositely charged ions (cations and anions)

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21
Q

What are the attractive forces associated with Covalent bonds

A

The attractive force between the nuclei of the bonding atoms and the shared bonding electrons

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22
Q

What are 3 similarities between ionic and covalent bonds?

A

Both form when atoms try to achieve a noble gas configuration

Both are strong when compared with intermolecular attractions

The energy when both types of bonds form is lower than the energy of the atoms apart

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23
Q

What are 3 differences between ionic and covalent bonds?

A

Ionic bonds form between metals and nonmetals, covalent bonds form between nonmetals

In ionic bonds there is a complete transfer of electrons, while in covalent bonds there is a sharing (equal or unequal) of electrons.

Compounds with ionic bonds are crystalline solids at room temperature while compounds with covalent bonds are solids, liquids, or gases at room temperature.

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24
Q

Glucose is a covalent compound with the molecular formula C6H12O6. This and many other covalent formulas are not reduced down to their simplest whole-number ratio of atoms in the compound. Explain why.

A

Glucose is a covalently bonded molecule composed of discrete molecules. Each molecule contains six carbon atoms, twelve hydrogen atoms and six oxygen atoms.
In contrast to ionic compounds, glucose does not form a crystal lattice

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25
Q

Many covalent compounds have much lower melting points than ionic compounds. Why doesn’t this mean that covalent bonds are weaker than ionic bonds?

A

Melting points of covalent compounds usually do not require breaking the bonds between atoms. Instead, intermolecular attractions are broken when melting covalent compounds. The actual bonds between the atoms do not get broken as they are quite strong

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26
Q

Diamond is a form of pure carbon containing only covalent bonds. It is the hardest substance known and a melting point of 3550o C. Explain its hardness and high melting point.

A

Diamonds are network covalent solids that are held together by covalent bonds that extend throughout the entire sample.

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27
Q

Consider the nature of the covalent bonds present in HCl and in N2. Which substance would you expect to have the higher melting point? Explain your answer

A

HCl would have the higher melting point. HCl is polar (ΔEN = 0.9) while N2 is non-polar.

Because HCl is polar it acts as a dipole, the H end is ∂+ while the Cl end is ∂-

∂+ end of one molecule lines up with the ∂- end of a different molecule setting up an electrostatic attraction which need to be broken in order to melt HCl.

N2 doesn’t have these attractions.

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28
Q

how to draw a dot structure

A

put number of valence electrons around element symbol

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29
Q

lewis structure

A

2 dimensional representation of the molecular formula, usually for covalent compounds

single lines represent bonds

other pairs of electrons are non-bonding or lone pairs

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30
Q

steps in drawing a Lewis structure

A
  1. determine total number of valence electrons
  2. draw the bonds between the atoms
  3. subtract the number of valence electrons used for bonding (each counts for 2 electrons)
  4. arrange remaining valence electrons to obey the octet rule
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31
Q

bond energy

A

energy required to break a mole of bonds

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32
Q

radical

A

A molecule with one or more unpaired electron in its outer shell

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33
Q

dimer

A

a molecule or molecular complex consisting of two identical molecules linked together

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34
Q

resonance structure

A

molecule or ion that contains double bonds next to single bonds, often has several possible structures

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35
Q

delocalized electrons

A

not associated with any one pair of bonded atoms. Are spread out equally between the three pairs of atoms

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36
Q

formal charge

A

charge that that an atom would have if all bonding electrons are shared equally between the bonded atoms

ignored electronegativity

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37
Q

formal charge =

A

number of valence electrons - (number of nonbonding electrons + 1/2 number of bonding electrons)

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38
Q

sum of formal charges in a neutral molecule

A

0

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39
Q

sum of formal charges in a polyatomic ion

A

charge of ion

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40
Q

small/zero formal charges on individual atoms are better than

A

large formal charges

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41
Q

when formal charge cannot be avoided, _ forlmal charge should reside on the most _ atom

A

negative, electronegative

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42
Q

what do the dots represent in a Lewis structure?

A

valence electrons

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43
Q

what do the elements symbol represent in a Lewis structure?

A

nucleus and core electrons

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44
Q

What do the number of dots in main group metals tell us about the charges of the ions formed by these metals?

A

magnitude of the positive charge

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45
Q

ABn

A

a is central atom bonded to n atoms of b

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46
Q

AB2

A

linear
180

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47
Q

AB3

A

trigonal planar/pyramidal
120

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48
Q

AB4

A

tetrahedral
109.5

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49
Q

AB5

A

trigonal bipyramidal
120
90

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50
Q

AB6

A

octahedral
90
120

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51
Q

VSEPR

A

explains molecular shapes for representative elements

Negatively charged electron domains repel each other

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52
Q

electron domains

A

electron pairs in a covalent bond

Note: Each multiple bond in a molecule also represents a single electron domain.

nonbonding pair of electrons

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53
Q

electron geometries

A

Best arrangement of electron domains minimizes repulsions among them.

Shapes of different ABn molecules or ions depend on number of electron domains surrounding the central atom.

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54
Q

molecular geometry

A

arrangement of only the atoms in a molecule or ion.

Nonbonding pairs are not part of the description.

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55
Q

steps to use VSEPR to predict molecular shapes

A
  1. draw Lewis structure
  2. determine electron domain geometry
  3. determine molecular geometry
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56
Q

linear molecular geometry

A

2 bonding domains
0 nonbonding domains

57
Q

trigonal planar molecular geometry

A

3 bonding domains
0 nonbonding domains

58
Q

bent molecular geometry

A

2 bonding domains
1/2 nonbonding domains

59
Q

tetrahedral molecular geometry

A

4 bonding domains
0 nonbonding domains

60
Q

trigonal pyramidal molecular geometry

A

3 bonding domains
1 nonbonding domains

61
Q

bond angles _ as nonbonding pairs increases

A

decrease

62
Q

electron domains for nonbonding pairs exert _ repulsive forces on adjacent electron domains

A

greater

63
Q

expanded valence shells are used when

A

there are 5 or 6 electron domains around the central atom

-central atom is in period 3 or above
-5 electron domains have one of four molecular geometries
-depends on number of nonbonding pairs and minimizing electron domain repulsions

64
Q

trigonal bipyramidal electron domain

A

2 axial domains

3 equatorial domains

each axial domain forms 90 degree angle with any equatorial domain

65
Q

seesaw electron domain

A

1 nonbonding domain

4 bonding

axial lone pair: 3 90 degree interaction with nonbonding pairs

equatorial lone pair: 2 90 degree interactions with bonding pairs

66
Q

t-shaped electron domains

A

2 nonbonding domains occupy 2 of 3 equatorial positions

3 bonding

67
Q

linear electron domains

A

3 nonbonding domains all occupy equatorial positions

68
Q

octahedral electron domains

A

0 nonbonding
6 bonding

69
Q

square pyramidal electron domain

A

1 nonbonding domain
4 bonding

70
Q

square planar electron domain

A

2 nonbonding
4 bonding

71
Q

bond polarity

A

measures how equally electrons in a bond are shared between 2 atoms of the bond

72
Q

bond polarity and electronegativity proportion

A

increases with the other

73
Q

dipole moment

A

measures the amount of charge separation in a diatomic molecule

74
Q

dipole moment of a non-diatomic molecule depends on

A

-polarities of individual bonds
-geometry of the molecule

75
Q

problems with Lewis Theory

A

-doesn’t give good :
-numerical predictions in property trends
-resonance predictions
-angle predictions
-correct magnetic behavior

76
Q

valence bond theory postulates

A

-buildup of electron density between 2 nuclei
-Overlap of valence atomic orbitals of two atoms.
-Always an optimum distance between two nuclei.
-valence electrons reside in quantum mechanical atomic orbitals (s,p,d,f)

77
Q

molecular orbitals

A

regions of high probability of finding shared electrons in the molecule

more stable than the separate atomic orbitals

78
Q

chemical bond results from

A

-the overlap of 2 half-filled orbitals with spin-pairing of the 2 valence electrons
-a completely filled orbital with an empty orbital

79
Q

geometry of overlapping orbitals determines

A

shape of molecule

80
Q

hybridizing

A

mixing different types of orbitals in the valence shell to make a new set of degenerate orbitals

molecules use the hybridization that yields the lowest overall energy for the molecule

81
Q

hybrid orbitals

A

have different shapes and energies from standard orbitals

still localized on individual atoms

82
Q

total number of standard atomic orbitals =

A

number of hybrid orbitals formed

83
Q

combinations of standard orbitals added determines

A

shapes and energies of hybrid orbitals

84
Q

pi bonds

A

p orbitals overlap side by side

results in electron density above and below internuclear axis

85
Q

sigma bonds

A

p orbitals that overlap end to end

86
Q

double bond

A

one sigma and one pi bond

87
Q

triple bond

A

one sigma and 2 pi bonds

88
Q

which are stronger? sigma or pi bonds

A

pi

89
Q

what types of bonds can you rotate around?

A

-single bonds are relatively unrestricted (if pi bond is broken)
-not double bonds

90
Q

isomerism

A

phenomenon in which more than one compounds have the same chemical formula but different chemical structures

91
Q

cis

A

same side

92
Q

trans

A

opposite side

93
Q

steps in predicting bonding in molecules

A
  1. draw lewis structure
  2. use VSEPR to predict electron geometry
94
Q

What is valence bond theory?

A

describes covalent bonding as a buildup ofelectron density between two nuclei when the valence atomic orbital of one atom overlaps with the valence atomic orbital of another atom

95
Q

What is a hybridized orbital?

A

A combination Q/ two standard (s, p, d or f) orbitals. Hybridized orbitals have different shapes from those of standard orbitals but are still localized on individual atoms

96
Q

Explain how the Lewis model and the valence bond theory differ in their description of a chemical bond

A

Lewis model: covalent bonds occur when atoms share electrons to concentrate electron density between the two nuclei.

Valence bond theory: a buildup of electron density when the valence atomic orbital of one atom overlaps the valence electron orbital of another atom

97
Q

In valence bond theory, what determines the shape of the molecule?

A

The combination of standard orbitals when added together

98
Q

bond order

A

number of bonding pairs of electrons between two of atoms and indicates the stability of a bond

99
Q

higher bond order =

A

greater stability of the molecule

more attraction between electrons (atoms held together more tightly)

100
Q

bond order and length proportion

A

inverse

101
Q

bond order and strength

A

direct

102
Q

to determine bond order for diatomic molecules:

A

draw Lewis structure and determine types of bonds between atoms

bond order 0: no bond
bond order 1: single
bond order 2: double
bond order 3: triple

103
Q

to determine bond order for non-diatomic molecules:

A
  1. draw Lewis structure
  2. count total number of bonds
  3. count number of bond groups between individual atoms
  4. divide number of bonds between atoms by the total number of bond groups in the molecule
104
Q

multivalent

A

atoms that can have more than one charge

105
Q

copper ions

A

1+ and 2+

106
Q

iron ions

A

2+ and 3+

107
Q

tin ions

A

2+ and 4+

108
Q

manganese ions

A

2+ and 4+

109
Q

lead ions

A

2+ and 4+

110
Q

silver ion

A

Ag1+

111
Q

Nickel ion

A

Ni2+

112
Q

zinc ion

A

Zn2+

113
Q

chromium

A

Cr3+

114
Q

ammonium

A

NH4(1+)

115
Q

hydrogen carbonate

A

HCO3(1-)

116
Q

chlorite

A

ClO2(1-)

117
Q

perchlorate

A

ClO4(1-)

118
Q

nitrate

A

NO3(1-)

119
Q

hydrogen sulfite

A

HSO3(1-)

120
Q

hydrogen sulfate

A

HSO4(1-)

121
Q

dichromate

A

Cr2O7(2-)

122
Q

thiosulfate

A

S2O3(2-)

123
Q

hydrogen phosphate

A

HPO4(2-)

124
Q

acetate

A

C2H3O2(1-)

125
Q

carbonate

A

CO3(2-)

126
Q

hypochlorite

A

ClO(1-)

127
Q

chlorate

A

ClO3(1-)

128
Q

nitrite

A

NO2(1-)

129
Q

sulfite

A

SO3(2-)

130
Q

sulfate

A

SO4(2-)

131
Q

permanganate

A

MnO4(1-)

132
Q

chromate

A

CrO2(2-)

133
Q

phosphate

A

PO4(3-)

134
Q

dihydrogen phosphate

A

H2PO4(2-)

135
Q

cyanide

A

CN(1-)

136
Q

oxyanions

A

an element bonded to oxygen atoms

137
Q

hydrates

A

ionic compounds that crystallize out of an aqueous solution and incorporate water molecules in a fixed ratio

138
Q

anhydrous

A

removing water from a hydrate by heating

139
Q
A

ap chem (11 decks)

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