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Bettelheim and Chang > Chemical Bonding > Flashcards

Chemical Bonding Flashcards

1
Q

atoms in compounds are held together by powerful forces of attraction called

A

Chemical bonds

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2
Q

Two main types of chemical bonds

A

Ionic and covalent bonds

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3
Q

devised a beautifully simple model that unified many of the observations about chemical bonding and chemical reactions

A

Gilbert N. Lewis

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4
Q

The tendency to react in ways that achieve an outer shell of eight valence electrons is particularly common among Group 1A–7A elements and is given the special name of the

A

octet rule

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5
Q

the atom becomes a negatively charged ion called an

A

anion

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6
Q

the atom becomes a positively charged ion called a

A

cation

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7
Q

When an ion forms, the number of ______ in the nucleus of the atom remains unchanged; only the number of ______ in the valence shell of the atom changes.

A

protons and neutrons; electrons

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8
Q

Imperfections of octet rule

A
  1. Ions of period 1 and 2 elements with charges greater than 12 are unstable.
  2. The octet rule does not apply to Group 1B–7B elements (the transition elements), most of which form ions with two or more different positive charges.
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9
Q

forms when a metal loses one or more valence electrons.

A

Monoatomic cation

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10
Q

the suffix ______ is used to show the smaller charge.

A

-ous

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11
Q

the suffix ______ is used to show the larger charge.

A

-ic

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12
Q

the name of the cation is the name of the metal followed by the word

A

ion

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13
Q

A monatomic anion is named by adding

A

-ide

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14
Q

contains more than one atom.

A

Polyatomic ion

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15
Q

A chemical bond resulting from the attraction between a positive ion and a negative ion

A

Ionic bond

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16
Q

A bond resulting from the sharing of electrons between two atoms.

A

covalent bond

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17
Q

atoms bond together in such a way that each atom participating in a bond acquires a valence-shell electron configuration matching that of the noble gas nearest to it in atomic number.

A

Lewis model of chemical bonding

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18
Q

is a measure of an atom’s attraction for the electrons it shares in a chemical bond with another atom.

A

Electronegativity

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19
Q

The most widely used scale of electronegativities was devised by

A

Linus Pauling

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20
Q

The more electronegative atom gains one or more valence electrons and becomes an

A

Anion

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21
Q

the less electronegative atom loses one or more valence electrons and becomes a

A

Cation

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22
Q

The compound formed by the combination of positive and negative ions is called an

A

Ionic compound

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23
Q

this type of electron transfer to form an ionic compound is most likely to occur if the difference in electronegativity between two atoms is approximately

A

1.9 or greater

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24
Q

contains only two elements

A

binary compound

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25
Q

both of the elements are present as ions.

A

binary ionic compound

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26
Q

A bond formed by sharing a pair of electrons is called a

A

single bond

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27
Q

Electronegativity less than 0.5

A

nonpolar covalent bond

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28
Q

Electronegativity 0.5-1.9

A

polar covalent bond

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29
Q

Greater than 1.9

A

ionic bond

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30
Q

electrons are shared equally

A

nonpolar covalent bond

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31
Q

they are shared unequally

A

polar covalent bond

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32
Q

A chemical species in which there is a separation of charge; there is a positive pole in one part of the species, and a negative pole in another part

A

dipole

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33
Q

the more electronegative atom gains a greater fraction of the shared electrons and acquires a

A

partial negative charge

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34
Q

The less electronegative atom has a lesser fraction of the shared electrons and acquires a

A

partial positive charge

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35
Q

A bond formed by sharing two pairs of electrons and represented by two lines between the two bonded atoms

A

double bond

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36
Q

A bond formed by sharing three pairs of electrons and represented by three lines between the two bonded atoms

A

triple bond

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37
Q

Exceptions to the octet rule

A

involves molecules that contain an atom with more than eight electrons in its valence shell.

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38
Q

is a binary (two-element) compound in which all bonds are covalent.

A

binary covalent compound

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39
Q

The theory of resonance, developed primarily by

A

Linus Pauling

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40
Q

A theory that many molecules and ions are best described as a hybrid of two or more Lewis contributing structures

A

Resonance

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41
Q

Representations of a molecule or ion that differ only in the distribution of valence electrons

A

contributing structure

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42
Q

A molecule or ion described as a composite or hybrid of a number of contributing structures

A

resonance hybrid

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43
Q

A symbol used to show that the structures on either side of it are resonance contributing structures

A

double headed arrow

44
Q

The angle between two atoms bonded to a central atom

A

bond angle

45
Q

We can predict bond angles in these and other molecules by using the

A

VSEPR (Valence shell electron pair repulsion)

46
Q

the valence electrons of an atom may be involved in the formation of single, double, or triple bonds, or they may be unshared.

A

VSEPR (Valence shell electron pair repulsion)

47
Q

The maximum separation occurs when the angle between any two regions of electron density is

A

109.5°

48
Q

all H C H bond angles to be 109.5°, and the shape of the molecule to be

A

tetrahedral

49
Q

the four regions are arranged in a tetrahedral manner and that the three H N H bond angles in this molecule are 109.5°. The observed bond angles are

A

107.3°

50
Q

The geometry of an ammonia molecule is described as

A

Pyramidal

51
Q

the molecule is shaped like a triangular-based pyramid with the three hydrogens located at the base and the single nitrogen located at the apex.

A

Pyramidal

52
Q

the actual H O H bond angle in a water molecule is

A

104.5°

53
Q

The geometry about an atom surrounded by three regions of electron density, as in formaldehyde and ethylene, is described as

A

trigonal planar

54
Q

A molecule will be polar if

A
  1. It has polar bonds
  2. Its centers of partial positive charge and partial negative charge lie at different places within the molecule.
55
Q

He said that atoms combine in order to achieve a more stable electron configuration

A

Gilbert Lewis

56
Q

consists of the symbol of an element and one dot for each valence electron in an atom of the element

A

Lewis dot symbol

57
Q

all have incompletely filled inner shells, and in general, we cannot write simple Lewis dot symbols for them

A

Transition metals, lanthanides, and actinides

58
Q

the electrostatic force that holds ions together in an ionic compound.

A

ionic bond

59
Q

the energy required to completely separate one mole of a solid ionic compound into gaseous ions

A

Lattice energy

60
Q

the potential energy (E) between two ions is directly proportional to the product of their charges and inversely proportional to the distance of separation between them

A

Coulomb’s Law

61
Q

relates lattice energies of ionic compounds to ionization energies, electron affinities, and other atomic and molecular properties

A

Bon Haber cycle

62
Q

The greater the lattice energy, the more stable the ionic compound

A

True

63
Q

a bond in which two electrons are shared by two atoms

A

Covalent bond

64
Q

compounds that contain only covalent bonds

A

Covalent compounds

65
Q

pairs of valence electrons that are not involved in covalent bond formation

A

lone pairs

66
Q

a representation of covalent bonding in which shared electron pairs are shown either as lines or as pairs of dots between two atoms, and lone pairs are shown as pairs of dots on individual atoms

A

Lewis structure

67
Q

An atom other than hydrogen tends to form bonds until it is surrounded by eight valence electrons

A

octet rule

68
Q

The octet rule works mainly for elements in the ______ of the periodic table

A

second period

69
Q

two atoms are held together by one electron pair

A

single bond

70
Q

two atoms share two or more pairs of electrons

A

multiple bonds

71
Q

two atoms share two pairs of electrons

A

double bonds

72
Q

two atoms share three pairs of electrons

A

triple bond

73
Q

distance between the nuclei of two covalently bonded atoms in a molecule

A

bond length

74
Q

two types of attractive forces in covalent compounds

A
  1. the force that holds the atoms together in a molecule
  2. intermolecular force
75
Q

the electrons spend more time in the vicinity of one atom than the other

A

polar covalent bonds

76
Q

sharing of electrons is exactly equal

A

nonpolar covalent bonds

77
Q

transfer of the electron(s) is nearly complete

A

ionic bond

78
Q

the ability of an atom to attract toward itself the electrons in a chemical bond

A

electronegativity

79
Q

the difference between the valence electrons in an isolated atom and the number of electrons assigned to that atom in a Lewis structure.

A

Formal charge

80
Q

one of two or more Lewis structures for a single molecule that cannot be represented accurately by only one Lewis structure.

A

Resonance structure

81
Q

use of two or more Lewis structures to represent a particular molecule

A

Resonance

82
Q

Exceptions to the octet rule: elements

A

Beryllium, Boron, Aluminum

83
Q

a covalent bond in which one of the atoms donates both electrons

A

coordinate covalent bond

84
Q

the enthalpy change required to break a particular bond in one mole of gaseous molecules

A

bond dissociation energy

85
Q

is the three-dimensional arrangement of atoms in a molecule

A

molecular geometry

86
Q

the outermost electron-occupied shell of an atom; it holds the electrons that are usually involved in bonding

A

valence shell

87
Q

responsible for holding two atoms together

A

bonding pair

88
Q

it accounts for the geometric arrangements of electron pairs around a central atom in terms of the electrostatic repulsion between electron pairs

A

VSEPR/ valence shell electron pair repulsion

89
Q

Two general rules about the use of VSEPR model:

A
  1. As far as electron-pair repulsion is concerned, double bonds and triple bonds can be treated like single bonds
  2. If a molecule has two or more resonance structures, we can apply the VSEPR model to any one of them.
90
Q

the product of the charge Q and the distance r between the charges

A

Dipole moment

91
Q

Diatomic molecules containing atoms of different elements have dipole moments and are called

A

Polar molecules

92
Q

Diatomic molecules containing atoms of the same element

A

Nonpolar molecules

93
Q

assumes that the electrons in a molecule occupy atomic orbitals of the individual atoms

A

valence bond theory

94
Q

assumes the formation of molecular orbitals from the atomic orbitals

A

molecular orbital theory

95
Q

atomic orbitals obtained when two or more nonequivalent orbitals of the same atom combine in preparation for covalent bond formation.

A

hybrid orbitals

96
Q

the mixing of atomic orbitals in an atom (usually a central atom) to generate a set of hybrid orbitals

A

hybridization

97
Q

covalent bonds formed by orbitals overlapping end-to-end, with the electron density concentrated between the nuclei of the bonding atoms

A

sigma bond

98
Q

a covalent bond formed by sideways overlapping orbitals with electron density concentrated above and below the plane of the nuclei of the bonding atoms.

A

pi bond

99
Q

has lower energy and greater stability than the atomic orbitals from which it was formed

A

bonding molecular orbital

100
Q

has higher energy and lower stability than the atomic orbitals from which it was formed

A

antibonding molecular orbital

101
Q

the electron density is concentrated symmetrically around a line between the two nuclei of the bonding atoms

A

sigma molecular orbital

102
Q

the electron density is concentrated above and below an imaginary line joining the two nuclei of the bonding atoms

A

pi molecular orbital

103
Q

indicates the strength of a bond

A

bond order

104
Q

diatomic molecules containing atoms of the same elements

A

homonuclear diatomic molecules

105
Q

are not confined between two adjacent bonding atoms, but actually extend over three or more atoms

A

delocalized molecular orbitals

Bettelheim and Chang (7 decks)

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