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unit 3 Flashcards

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
Many covalent compounds have much lower melting points than ionic compounds. Why doesn’t this mean that covalent bonds are weaker than ionic bonds?
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
26
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.
Diamonds are network covalent solids that are held together by covalent bonds that extend throughout the entire sample.
27
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
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.
28
how to draw a dot structure
put number of valence electrons around element symbol
29
lewis structure
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
30
steps in drawing a Lewis structure
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
31
bond energy
energy required to break a mole of bonds
32
radical
A molecule with one or more unpaired electron in its outer shell
33
dimer
a molecule or molecular complex consisting of two identical molecules linked together
34
resonance structure
molecule or ion that contains double bonds next to single bonds, often has several possible structures
35
delocalized electrons
not associated with any one pair of bonded atoms. Are spread out equally between the three pairs of atoms
36
formal charge
charge that that an atom would have if all bonding electrons are shared equally between the bonded atoms ignored electronegativity
37
formal charge =
number of valence electrons - (number of nonbonding electrons + 1/2 number of bonding electrons)
38
sum of formal charges in a neutral molecule
0
39
sum of formal charges in a polyatomic ion
charge of ion
40
small/zero formal charges on individual atoms are better than
large formal charges
41
when formal charge cannot be avoided, _ forlmal charge should reside on the most _ atom
negative, electronegative
42
what do the dots represent in a Lewis structure?
valence electrons
43
what do the elements symbol represent in a Lewis structure?
nucleus and core electrons
44
What do the number of dots in main group metals tell us about the charges of the ions formed by these metals?
magnitude of the positive charge
45
ABn
a is central atom bonded to n atoms of b
46
AB2
linear 180
47
AB3
trigonal planar/pyramidal 120
48
AB4
tetrahedral 109.5
49
AB5
trigonal bipyramidal 120 90
50
AB6
octahedral 90 120
51
VSEPR
explains molecular shapes for representative elements Negatively charged electron domains repel each other
52
electron domains
electron pairs in a covalent bond Note: Each multiple bond in a molecule also represents a single electron domain. nonbonding pair of electrons
53
electron geometries
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.
54
molecular geometry
arrangement of only the atoms in a molecule or ion. Nonbonding pairs are not part of the description.
55
steps to use VSEPR to predict molecular shapes
1. draw Lewis structure 2. determine electron domain geometry 3. determine molecular geometry
56
linear molecular geometry
2 bonding domains 0 nonbonding domains
57
trigonal planar molecular geometry
3 bonding domains 0 nonbonding domains
58
bent molecular geometry
2 bonding domains 1/2 nonbonding domains
59
tetrahedral molecular geometry
4 bonding domains 0 nonbonding domains
60
trigonal pyramidal molecular geometry
3 bonding domains 1 nonbonding domains
61
bond angles _ as nonbonding pairs increases
decrease
62
electron domains for nonbonding pairs exert _ repulsive forces on adjacent electron domains
greater
63
expanded valence shells are used when
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
trigonal bipyramidal electron domain
2 axial domains 3 equatorial domains each axial domain forms 90 degree angle with any equatorial domain
65
seesaw electron domain
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
t-shaped electron domains
2 nonbonding domains occupy 2 of 3 equatorial positions 3 bonding
67
linear electron domains
3 nonbonding domains all occupy equatorial positions
68
octahedral electron domains
0 nonbonding 6 bonding
69
square pyramidal electron domain
1 nonbonding domain 4 bonding
70
square planar electron domain
2 nonbonding 4 bonding
71
bond polarity
measures how equally electrons in a bond are shared between 2 atoms of the bond
72
bond polarity and electronegativity proportion
increases with the other
73
dipole moment
measures the amount of charge separation in a diatomic molecule
74
dipole moment of a non-diatomic molecule depends on
-polarities of individual bonds -geometry of the molecule
75
problems with Lewis Theory
-doesn't give good : -numerical predictions in property trends -resonance predictions -angle predictions -correct magnetic behavior
76
valence bond theory postulates
-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
molecular orbitals
regions of high probability of finding shared electrons in the molecule more stable than the separate atomic orbitals
78
chemical bond results from
-the overlap of 2 half-filled orbitals with spin-pairing of the 2 valence electrons -a completely filled orbital with an empty orbital
79
geometry of overlapping orbitals determines
shape of molecule
80
hybridizing
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
hybrid orbitals
have different shapes and energies from standard orbitals still localized on individual atoms
82
total number of standard atomic orbitals =
number of hybrid orbitals formed
83
combinations of standard orbitals added determines
shapes and energies of hybrid orbitals
84
pi bonds
p orbitals overlap side by side results in electron density above and below internuclear axis
85
sigma bonds
p orbitals that overlap end to end
86
double bond
one sigma and one pi bond
87
triple bond
one sigma and 2 pi bonds
88
which are stronger? sigma or pi bonds
pi
89
what types of bonds can you rotate around?
-single bonds are relatively unrestricted (if pi bond is broken) -not double bonds
90
isomerism
phenomenon in which more than one compounds have the same chemical formula but different chemical structures
91
cis
same side
92
trans
opposite side
93
steps in predicting bonding in molecules
1. draw lewis structure 2. use VSEPR to predict electron geometry
94
What is valence bond theory?
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
What is a hybridized orbital?
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
Explain how the Lewis model and the valence bond theory differ in their description of a chemical bond
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
In valence bond theory, what determines the shape of the molecule?
The combination of standard orbitals when added together
98
bond order
number of bonding pairs of electrons between two of atoms and indicates the stability of a bond
99
higher bond order =
greater stability of the molecule more attraction between electrons (atoms held together more tightly)
100
bond order and length proportion
inverse
101
bond order and strength
direct
102
to determine bond order for diatomic molecules:
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
to determine bond order for non-diatomic molecules:
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
multivalent
atoms that can have more than one charge
105
copper ions
1+ and 2+
106
iron ions
2+ and 3+
107
tin ions
2+ and 4+
108
manganese ions
2+ and 4+
109
lead ions
2+ and 4+
110
silver ion
Ag1+
111
Nickel ion
Ni2+
112
zinc ion
Zn2+
113
chromium
Cr3+
114
ammonium
NH4(1+)
115
hydrogen carbonate
HCO3(1-)
116
chlorite
ClO2(1-)
117
perchlorate
ClO4(1-)
118
nitrate
NO3(1-)
119
hydrogen sulfite
HSO3(1-)
120
hydrogen sulfate
HSO4(1-)
121
dichromate
Cr2O7(2-)
122
thiosulfate
S2O3(2-)
123
hydrogen phosphate
HPO4(2-)
124
acetate
C2H3O2(1-)
125
carbonate
CO3(2-)
126
hypochlorite
ClO(1-)
127
chlorate
ClO3(1-)
128
nitrite
NO2(1-)
129
sulfite
SO3(2-)
130
sulfate
SO4(2-)
131
permanganate
MnO4(1-)
132
chromate
CrO2(2-)
133
phosphate
PO4(3-)
134
dihydrogen phosphate
H2PO4(2-)
135
cyanide
CN(1-)
136
oxyanions
an element bonded to oxygen atoms
137
hydrates
ionic compounds that crystallize out of an aqueous solution and incorporate water molecules in a fixed ratio
138
anhydrous
removing water from a hydrate by heating
139

ap chem (11 decks)

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