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IMAT Chem > Chapter 3: Chemical Bonding > Flashcards

Chapter 3: Chemical Bonding Flashcards

1
Q

What is ionic bonding?

A

Electrostatic attraction between positive and negative ions.

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2
Q

How are positive ions formed?

A

By losing one or more electrons.

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3
Q

How are negative ions formed?

A

By gaining one or more electrons.

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4
Q

Why do elements in Group 14 generally not form ions?

A

They have four valence electrons and no tendency to gain or lose electrons.

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5
Q

What is the structure of ionic compounds?

A

A lattice structure.

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6
Q

What is a key property of ionic compounds related to temperature?

A

High melting and boiling points.

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7
Q

In what state do ionic compounds conduct electricity?

A

Molten or dissolved in water

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8
Q

What is a covalent bond?

A

A bond formed by sharing valence electrons.

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9
Q

What determines bond polarity?

A

The difference in electronegativity between bonded atoms.

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10
Q

How does bond strength change with bond length?

A

Shorter bonds are stronger.

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11
Q

What is the octet rule?

A

Atoms tend to have eight electrons in their valence shell.

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12
Q

Which elements often form stable compounds with incomplete octets?

A

Beryllium (Be) and Boron (B).

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13
Q

What are resonance structures?

A

Multiple valid arrangements of electrons in a molecule.

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14
Q

What theory determines molecular shape?

A

VSEPR theory (Valence Shell Electron Pair Repulsion).

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15
Q

How does electron domain affect molecular geometry?

A

Electron domains repel and spread as far as possible.

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16
Q

What is a lone pair?

A

A pair of valence electrons not involved in bonding.

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17
Q

What is an expanded octet?

A

More than 8 electrons around a central atom.

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18
Q

Why can atoms have expanded octets?

A

d orbitals in valence shell with energy close to p orbitals are available. Electron from p promote to occupy d.

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19
Q

Which elements can have expanded octets?

A

Period 3 and below (e.g. Cl, S, F, etc.)

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20
Q

What are delocalized electrons?

A

Electrons free to move in a metallic lattice.

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21
Q

What increases the strength of metallic bonding?

A

Higher positive charge

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22
Q

What type of molecules experience London dispersion forces?

A

All molecules

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23
Q

What is dipole-dipole attraction?

A

Attraction between permanent dipoles in polar molecules.

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24
Q

What is hydrogen bonding?

A

A strong dipole-dipole attraction involving H and highly electronegative atoms (N, O, F,..)

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25
Why do substances with hydrogen bonds have high boiling points?
Hydrogen bonding is the strongest intermolecular force.
26
What are giant covalent substances?
Substances with continuous covalent networks.
27
What is the electron geometry for 2 electron domains?
Linear.
28
What is the electron geometry for 3 electron domains?
Trigonal planar.
29
What is the electron geometry for 4 electron domains?
Tetrahedral.
30
Why is H2O bent but BeH2 linear?
Lone pairs on oxygen in H2O cause repulsion.
31
Why is NH3's bond angle smaller than CH4's?
Lone pairs on nitrogen cause greater repulsion.
32
What is the molecular geometry of CF4?
Tetrahedral.
33
What is the molecular geometry of BF3?
Trigonal planar.
34
Why do ionic compounds dissolve in polar solvents?
Attraction between ions and solvent molecules.
35
What is the primary feature of a metallic lattice?
A regular arrangement of metal cations surrounded by a sea of delocalized electrons.
36
Which molecules have higher boiling points: polar or nonpolar?
Polar molecules.
37
What causes the bending shape in molecules like water?
Lone pairs causing greater electron repulsion.
38
What defines molecular polarity?
Shape and arrangement of polar bonds.
39
Why do metals conduct electricity?
Free-moving delocalized electrons.
40
What is a key feature of covalent bonds in terms of shared electrons?
Single bonds share 2 electrons; double share 4
41
What is the weakest intermolecular force?
London dispersion forces.
42
Why do nonpolar molecules generally have low boiling points?
Weak London dispersion forces => little energy needed to break the bonds.
43
How is bond energy related to bond strength?
Higher bond energy indicates stronger bonds.
44
What does VSEPR stand for?
Valence Shell Electron Pair Repulsion.
45
Why are metals malleable?
Layers of cations can slide over each other.
46
What is the significance of electron domain geometry?
Determines the spatial arrangement of electron pairs.
47
What does the term 'ionic lattice' describe?
A 3D repeating arrangement of ions.
48
Why are ionic compounds brittle?
Shifting layers cause like charges to repel when affected by force.
49
Why does Cl2 have a low boiling point?
Weak London dispersion forces between nonpolar molecules.
50
What is the shape of SO2?
Bent.
51
Why does metallic bonding increase with more delocalized electrons?
More electrons contribute to the attraction.
52
What causes dipole-dipole attractions?
Uneven electron distribution creating partial charges .
53
How does molecular size affect London dispersion forces?
Larger molecules have stronger dispersion forces.
54
What is the shape of PCl3?
Trigonal pyramidal.
55
What are the properties of giant covalent substances?
High melting points, non-soluble in all solutions, non-conductors | except *graphite, graphene*, and semi-conductivity of *Si*, *fullerene*
56
Why does NH4+ have identical bond angles to CH4?
Symmetrical distribution of bonding pairs.
57
What is the molecular geometry of Cl2CO?
Trigonal planar.
58
What determines a molecule's polarity?
Bond polarity and molecular geometry.
59
What is the general trend for melting points of ionic compounds?
High melting points due to strong electrostatic forces.
60
What is the structure of BeH2?
Linear.
61
Why is CH3+ an incomplete octet?
Carbon has only 6 valence electrons in this ion.
62
What defines bond strength in covalent bonds?
Number of shared electrons.
63
Why does CO2 have a linear shape?
Two double bonds and no lone pairs on the central atom.
64
What are delocalized electrons responsible for in metals?
Electrical and thermal conductivity.
65
Why do polar substances dissolve in water?
Water's polarity interacts with solute dipoles.
66
What is the shape of NH3?
Trigonal pyramidal.
67
What are the characteristics of polar covalent bonds?
Unequal sharing of electrons.
68
Why does SF6 have an expanded octet?
Sulfur uses d orbitals to accommodate extra electrons.
69
What is the molecular shape of NO2-?
Bent.
70
What makes metals ductile?
Ability of layers to slide under force.
71
Why do noble gases have low boiling points?
Weak London dispersion forces.
72
What increases the strength of dipole-dipole interactions?
Greater polarity.
73
Why is SO2F2 polar?
Asymmetrical shape with polar bonds.
74
What is metallic bonding?
Electrostatic attraction between cations and delocalized electrons.
75
What happens to metal atoms in a lattice?
They lose outer electrons and become cations.
76
What are the three main intermolecular forces?
London dispersion forces, dipole-dipole attraction, Hydrogen bonds | Van der Waals Force
77
What are London dispersion forces caused by?
Temporary dipoles inducing neighboring molecules, creating other temporary dipoles.
78
How do London dispersion forces change with molecular size?
They increase with larger molecular size.
79
Which substances only experience London dispersion forces?
Nonpolar molecules.
80
What effect does hydrogen bonding have on boiling points?
It significantly increases boiling points.
81
What is the strongest type of intermolecular force?
Hydrogen bonding.
82
Which atoms can form hydrogen bonds with hydrogen?
Oxygen (O), Fluorine (F), Nitrogen (N),...
83
What are the physical states of many nonpolar molecules at room temperature?
Gases.
84
Why do polar molecules have higher melting points than nonpolar ones? | between substances with comparable molecular mass
Stronger dipole-dipole attractions.
85
What causes a molecule to be nonpolar despite containing polar bonds?
*Symmetrical* arrangement of bonds cancels dipoles **and** bonds are of *equal* polarity
86
What is the molecular geometry of ClO2-?
Bent.
87
What is the polarity of CF4?
Nonpolar
88
What is the molecular shape of CS2?
Linear.
89
What is the difference between electron-pair geometry and molecular shape?
Electron-pair geometry considers all electron domains; Molecular shape considers only bonding pairs, not lone pairs.
90
What is the molecular structure of ClOF2+?
T-shaped.
91
Why does the bond angle in H-O-H differ from H-Be-H?
Lone pairs on oxygen in H-O-H cause more repulsion than bonded pairs.
92
What defines molecular geometry of a species with 3 bonding domains?
Trigonal planar geometry.
93
What is the molecular geometry of SO2F2?
Tetrahedral.
94
Why do covalent substances generally have low boiling points?
Weak intermolecular forces compared to ionic bonds.
95
What is the polarity of NH3?
Polar
96
What happens to boiling points as hydrogen bonding increases?
Boiling points rise significantly.
97
What is the general trend in the strength of intermolecular forces?
London dispersion < dipole-dipole < hydrogen bonding.
98
What causes volatility in covalent substances?
Weak intermolecular forces.
99
Why are ionic compounds often brittle?
When layers of ions shift, ions with same charges meet and repel each other => break structure
100
What type of substances dissolve in polar solvents?
Polar and ionic compounds.

IMAT Chem (9 decks)

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