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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
Q

Why do substances with hydrogen bonds have high boiling points?

A

Hydrogen bonding is the strongest intermolecular force.

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26
Q

What are giant covalent substances?

A

Substances with continuous covalent networks.

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27
Q

What is the electron geometry for 2 electron domains?

A

Linear.

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28
Q

What is the electron geometry for 3 electron domains?

A

Trigonal planar.

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29
Q

What is the electron geometry for 4 electron domains?

A

Tetrahedral.

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30
Q

Why is H2O bent but BeH2 linear?

A

Lone pairs on oxygen in H2O cause repulsion.

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31
Q

Why is NH3’s bond angle smaller than CH4’s?

A

Lone pairs on nitrogen cause greater repulsion.

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32
Q

What is the molecular geometry of CF4?

A

Tetrahedral.

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33
Q

What is the molecular geometry of BF3?

A

Trigonal planar.

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34
Q

Why do ionic compounds dissolve in polar solvents?

A

Attraction between ions and solvent molecules.

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35
Q

What is the primary feature of a metallic lattice?

A

A regular arrangement of metal cations surrounded by a sea of delocalized electrons.

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36
Q

Which molecules have higher boiling points: polar or nonpolar?

A

Polar molecules.

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37
Q

What causes the bending shape in molecules like water?

A

Lone pairs causing greater electron repulsion.

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38
Q

What defines molecular polarity?

A

Shape and arrangement of polar bonds.

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39
Q

Why do metals conduct electricity?

A

Free-moving delocalized electrons.

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40
Q

What is a key feature of covalent bonds in terms of shared electrons?

A

Single bonds share 2 electrons; double share 4

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41
Q

What is the weakest intermolecular force?

A

London dispersion forces.

42
Q

Why do nonpolar molecules generally have low boiling points?

A

Weak London dispersion forces => little energy needed to break the bonds.

43
Q

How is bond energy related to bond strength?

A

Higher bond energy indicates stronger bonds.

44
Q

What does VSEPR stand for?

A

Valence Shell Electron Pair Repulsion.

45
Q

Why are metals malleable?

A

Layers of cations can slide over each other.

46
Q

What is the significance of electron domain geometry?

A

Determines the spatial arrangement of electron pairs.

47
Q

What does the term ‘ionic lattice’ describe?

A

A 3D repeating arrangement of ions.

48
Q

Why are ionic compounds brittle?

A

Shifting layers cause like charges to repel when affected by force.

49
Q

Why does Cl2 have a low boiling point?

A

Weak London dispersion forces between nonpolar molecules.

50
Q

What is the shape of SO2?

A

Bent.

51
Q

Why does metallic bonding increase with more delocalized electrons?

A

More electrons contribute to the attraction.

52
Q

What causes dipole-dipole attractions?

A

Uneven electron distribution creating partial charges .

53
Q

How does molecular size affect London dispersion forces?

A

Larger molecules have stronger dispersion forces.

54
Q

What is the shape of PCl3?

A

Trigonal pyramidal.

55
Q

What are the properties of giant covalent substances?

A

High melting points, non-soluble in all solutions, non-conductors

except graphite, graphene, and semi-conductivity of Si, fullerene

56
Q

Why does NH4+ have identical bond angles to CH4?

A

Symmetrical distribution of bonding pairs.

57
Q

What is the molecular geometry of Cl2CO?

A

Trigonal planar.

58
Q

What determines a molecule’s polarity?

A

Bond polarity and molecular geometry.

59
Q

What is the general trend for melting points of ionic compounds?

A

High melting points due to strong electrostatic forces.

60
Q

What is the structure of BeH2?

A

Linear.

61
Q

Why is CH3+ an incomplete octet?

A

Carbon has only 6 valence electrons in this ion.

62
Q

What defines bond strength in covalent bonds?

A

Number of shared electrons.

63
Q

Why does CO2 have a linear shape?

A

Two double bonds and no lone pairs on the central atom.

64
Q

What are delocalized electrons responsible for in metals?

A

Electrical and thermal conductivity.

65
Q

Why do polar substances dissolve in water?

A

Water’s polarity interacts with solute dipoles.

66
Q

What is the shape of NH3?

A

Trigonal pyramidal.

67
Q

What are the characteristics of polar covalent bonds?

A

Unequal sharing of electrons.

68
Q

Why does SF6 have an expanded octet?

A

Sulfur uses d orbitals to accommodate extra electrons.

69
Q

What is the molecular shape of NO2-?

A

Bent.

70
Q

What makes metals ductile?

A

Ability of layers to slide under force.

71
Q

Why do noble gases have low boiling points?

A

Weak London dispersion forces.

72
Q

What increases the strength of dipole-dipole interactions?

A

Greater polarity.

73
Q

Why is SO2F2 polar?

A

Asymmetrical shape with polar bonds.

74
Q

What is metallic bonding?

A

Electrostatic attraction between cations and delocalized electrons.

75
Q

What happens to metal atoms in a lattice?

A

They lose outer electrons and become cations.

76
Q

What are the three main intermolecular forces?

A

London dispersion forces, dipole-dipole attraction, Hydrogen bonds

Van der Waals Force

77
Q

What are London dispersion forces caused by?

A

Temporary dipoles inducing neighboring molecules, creating other temporary dipoles.

78
Q

How do London dispersion forces change with molecular size?

A

They increase with larger molecular size.

79
Q

Which substances only experience London dispersion forces?

A

Nonpolar molecules.

80
Q

What effect does hydrogen bonding have on boiling points?

A

It significantly increases boiling points.

81
Q

What is the strongest type of intermolecular force?

A

Hydrogen bonding.

82
Q

Which atoms can form hydrogen bonds with hydrogen?

A

Oxygen (O), Fluorine (F), Nitrogen (N),…

83
Q

What are the physical states of many nonpolar molecules at room temperature?

A

Gases.

84
Q

Why do polar molecules have higher melting points than nonpolar ones?

between substances with comparable molecular mass

A

Stronger dipole-dipole attractions.

85
Q

What causes a molecule to be nonpolar despite containing polar bonds?

A

Symmetrical arrangement of bonds cancels dipoles and bonds are of equal polarity

86
Q

What is the molecular geometry of ClO2-?

A

Bent.

87
Q

What is the polarity of CF4?

A

Nonpolar

88
Q

What is the molecular shape of CS2?

A

Linear.

89
Q

What is the difference between electron-pair geometry and molecular shape?

A

Electron-pair geometry considers all electron domains;

Molecular shape considers only bonding pairs, not lone pairs.

90
Q

What is the molecular structure of ClOF2+?

A

T-shaped.

91
Q

Why does the bond angle in H-O-H differ from H-Be-H?

A

Lone pairs on oxygen in H-O-H cause more repulsion than bonded pairs.

92
Q

What defines molecular geometry of a species with 3 bonding domains?

A

Trigonal planar geometry.

93
Q

What is the molecular geometry of SO2F2?

A

Tetrahedral.

94
Q

Why do covalent substances generally have low boiling points?

A

Weak intermolecular forces compared to ionic bonds.

95
Q

What is the polarity of NH3?

A

Polar

96
Q

What happens to boiling points as hydrogen bonding increases?

A

Boiling points rise significantly.

97
Q

What is the general trend in the strength of intermolecular forces?

A

London dispersion < dipole-dipole < hydrogen bonding.

98
Q

What causes volatility in covalent substances?

A

Weak intermolecular forces.

99
Q

Why are ionic compounds often brittle?

A

When layers of ions shift, ions with same charges meet and repel each other
=> break structure

100
Q

What type of substances dissolve in polar solvents?

A

Polar and ionic compounds.

IMAT Chem (9 decks)

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