Chemical Bonding Flashcards
1
Q
atoms in compounds are held together by powerful forces of attraction called
A
Chemical bonds
2
Q
Two main types of chemical bonds
A
Ionic and covalent bonds
3
Q
devised a beautifully simple model that unified many of the observations about chemical bonding and chemical reactions
A
Gilbert N. Lewis
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
5
Q
the atom becomes a negatively charged ion called an
A
anion
6
Q
the atom becomes a positively charged ion called a
A
cation
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
8
Q
Imperfections of octet rule
A
- Ions of period 1 and 2 elements with charges greater than 12 are unstable.
- 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.
9
Q
forms when a metal loses one or more valence electrons.
A
Monoatomic cation
10
Q
the suffix ______ is used to show the smaller charge.
A
-ous
11
Q
the suffix ______ is used to show the larger charge.
A
-ic
12
Q
the name of the cation is the name of the metal followed by the word
A
ion
13
Q
A monatomic anion is named by adding
A
-ide
14
Q
contains more than one atom.
A
Polyatomic ion
15
Q
A chemical bond resulting from the attraction between a positive ion and a negative ion
A
Ionic bond
16
Q
A bond resulting from the sharing of electrons between two atoms.
A
covalent bond
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
18
Q
is a measure of an atom’s attraction for the electrons it shares in a chemical bond with another atom.
A
Electronegativity
19
Q
The most widely used scale of electronegativities was devised by
A
Linus Pauling
20
Q
The more electronegative atom gains one or more valence electrons and becomes an
A
Anion
21
Q
the less electronegative atom loses one or more valence electrons and becomes a
A
Cation
22
Q
The compound formed by the combination of positive and negative ions is called an
A
Ionic compound
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
24
Q
contains only two elements
A
binary compound
25
both of the elements are present as ions.
binary ionic compound
26
A bond formed by sharing a pair of electrons is called a
single bond
27
Electronegativity less than 0.5
nonpolar covalent bond
28
Electronegativity 0.5-1.9
polar covalent bond
29
Greater than 1.9
ionic bond
30
electrons are shared equally
nonpolar covalent bond
31
they are shared unequally
polar covalent bond
32
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
dipole
33
the more electronegative atom gains a greater fraction of the shared electrons and acquires a
partial negative charge
34
The less electronegative atom has a lesser fraction of the shared electrons and acquires a
partial positive charge
35
A bond formed by sharing two pairs of electrons and represented by two lines between the two bonded atoms
double bond
36
A bond formed by sharing three pairs of electrons and represented by three lines between the two bonded atoms
triple bond
37
Exceptions to the octet rule
involves molecules that contain an atom with more than eight electrons in its valence shell.
38
is a binary (two-element) compound in which all bonds are covalent.
binary covalent compound
39
The theory of resonance, developed primarily by
Linus Pauling
40
A theory that many molecules and ions are best described as a hybrid of two or more Lewis contributing structures
Resonance
41
Representations of a molecule or ion that differ only in the distribution of valence electrons
contributing structure
42
A molecule or ion described as a composite or hybrid of a number of contributing structures
resonance hybrid
43
A symbol used to show that the structures on either side of it are resonance contributing structures
double headed arrow
44
The angle between two atoms bonded to a central atom
bond angle
45
We can predict bond angles in these and other molecules by using the
VSEPR (Valence shell electron pair repulsion)
46
the valence electrons of an atom may be involved in the formation of single, double, or triple bonds, or they may be unshared.
VSEPR (Valence shell electron pair repulsion)
47
The maximum separation occurs when the angle between any two regions of electron density is
109.5°
48
all H C H bond angles to be 109.5°, and the shape of the molecule to be
tetrahedral
49
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
107.3°
50
The geometry of an ammonia molecule is described as
Pyramidal
51
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.
Pyramidal
52
the actual H O H bond angle in a water molecule is
104.5°
53
The geometry about an atom surrounded by three regions of electron density, as in formaldehyde and ethylene, is described as
trigonal planar
54
A molecule will be polar if
1. It has polar bonds
2. Its centers of partial positive charge and partial negative charge lie at different places within the molecule.
55
He said that atoms combine in order to achieve a more stable electron configuration
Gilbert Lewis
56
consists of the symbol of an element and one dot for each valence electron in an atom of the element
Lewis dot symbol
57
all have incompletely filled inner shells, and in general, we cannot write simple Lewis dot symbols for them
Transition metals, lanthanides, and actinides
58
the electrostatic force that holds ions together in an ionic compound.
ionic bond
59
the energy required to completely separate one mole of a solid ionic compound into gaseous ions
Lattice energy
60
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
Coulomb's Law
61
relates lattice energies of ionic compounds to ionization energies, electron affinities, and other atomic and molecular properties
Bon Haber cycle
62
The greater the lattice energy, the more stable the ionic compound
True
63
a bond in which two electrons are shared by two atoms
Covalent bond
64
compounds that contain only covalent bonds
Covalent compounds
65
pairs of valence electrons that are not involved in covalent bond formation
lone pairs
66
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
Lewis structure
67
An atom other than hydrogen tends to form bonds until it is surrounded by eight valence electrons
octet rule
68
The octet rule works mainly for elements in the ______ of the periodic table
second period
69
two atoms are held together by one electron pair
single bond
70
two atoms share two or more pairs of electrons
multiple bonds
71
two atoms share two pairs of electrons
double bonds
72
two atoms share three pairs of electrons
triple bond
73
distance between the nuclei of two covalently bonded atoms in a molecule
bond length
74
two types of attractive forces in covalent compounds
1. the force that holds the atoms together in a molecule
2. intermolecular force
75
the electrons spend more time in the vicinity of one atom than the other
polar covalent bonds
76
sharing of electrons is exactly equal
nonpolar covalent bonds
77
transfer of the electron(s) is nearly complete
ionic bond
78
the ability of an atom to attract toward itself the electrons in a chemical bond
electronegativity
79
the difference between the valence electrons in an isolated atom and the number of electrons assigned to that atom in a Lewis structure.
Formal charge
80
one of two or more Lewis structures for a single molecule that cannot be represented accurately by only one Lewis structure.
Resonance structure
81
use of two or more Lewis structures to represent a particular molecule
Resonance
82
Exceptions to the octet rule: elements
Beryllium, Boron, Aluminum
83
a covalent bond in which one of the atoms donates both electrons
coordinate covalent bond
84
the enthalpy change required to break a particular bond in one mole of gaseous molecules
bond dissociation energy
85
is the three-dimensional arrangement of atoms in a molecule
molecular geometry
86
the outermost electron-occupied shell of an atom; it holds the electrons that are usually involved in bonding
valence shell
87
responsible for holding two atoms together
bonding pair
88
it accounts for the geometric arrangements of electron pairs around a central atom in terms of the electrostatic repulsion between electron pairs
VSEPR/ valence shell electron pair repulsion
89
Two general rules about the use of VSEPR model:
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
the product of the charge Q and the distance r between the charges
Dipole moment
91
Diatomic molecules containing atoms of different elements have dipole moments and are called
Polar molecules
92
Diatomic molecules containing atoms of the same element
Nonpolar molecules
93
assumes that the electrons in a molecule occupy atomic orbitals of the individual atoms
valence bond theory
94
assumes the formation of molecular orbitals from the atomic orbitals
molecular orbital theory
95
atomic orbitals obtained when two or more nonequivalent orbitals of the same atom combine in preparation for covalent bond formation.
hybrid orbitals
96
the mixing of atomic orbitals in an atom (usually a central atom) to generate a set of hybrid orbitals
hybridization
97
covalent bonds formed by orbitals overlapping end-to-end, with the electron density concentrated between the nuclei of the bonding atoms
sigma bond
98
a covalent bond formed by sideways overlapping orbitals with electron density concentrated above and below the plane of the nuclei of the bonding atoms.
pi bond
99
has lower energy and greater stability than the atomic orbitals from which it was formed
bonding molecular orbital
100
has higher energy and lower stability than the atomic orbitals from which it was formed
antibonding molecular orbital
101
the electron density is concentrated symmetrically around a line between the two nuclei of the bonding atoms
sigma molecular orbital
102
the electron density is concentrated above and below an imaginary line joining the two nuclei of the bonding atoms
pi molecular orbital
103
indicates the strength of a bond
bond order
104
diatomic molecules containing atoms of the same elements
homonuclear diatomic molecules
105
are not confined between two adjacent bonding atoms, but actually extend over three or more atoms
delocalized molecular orbitals
