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×Chemistry 1 > Chemical Bonding > Flashcards

Chemical Bonding Flashcards

1
Q

chemical bond

A

mutual electrical attraction between the nuclei and valence electrons of different atoms that binds the atoms together

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2
Q

purpose of chemical bonds

A

decrease potential energy of atoms

increase stability of atoms

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3
Q

types of chemical bonds

A

ionic bond

covalent bond

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4
Q

difference between types of chemical bonds

A

the way the electrons are redistributed

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5
Q

ionic bond

A

the transfer of electrons to other atoms

results from the electrical attraction between large numbers of cations and anions

bond between atoms with an electronegativity difference of 1.7 or greater

formed between metals and nonmetals

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6
Q

covalent bond

A

the sharing of electrons between two bonded atoms

bond between atoms with an electronegativity difference of 1.7 or less

formed between nonmetals

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7
Q

determining factor of type of chemical bond

A

difference in the bonded atoms’ electronegativities

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8
Q

type of bond between atoms of the same element

A

covalent bond

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9
Q

types of covalent bonds

A

nonpolar covalent bond

polar covalent bond

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10
Q

nonpolar covalent bond

A

covalent bond in which the bonding electrons are shared equally by the bonded atoms

results in a balanced distribution of electrical charge

bond between atoms with an electronegativity difference between 0.0 and 0.3

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11
Q

polar covalent bond

A

covalent bond in which the bonded atoms have an unequal attraction for the shared electrons

results in an imbalanced distribution of electrical charge

bond between atoms with an electronegativity difference between 0.3 and 1.7

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12
Q

molecule

A

a neutral group of atoms that are held together by covalent bonds

simplest unit of a molecular compound

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13
Q

molecular compound

covalent compound

A

chemical compound formed by two or more bonded atoms through the sharing of electrons

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14
Q

chemical formula

A

indication of the relative numbers of atoms of each kind in a chemical compound by using atomic symbols and numerical subscripts

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15
Q

molecular formula

A

chemical formula of a molecular compound

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16
Q

diatomic molecule

A

a molecule containing only two atoms

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17
Q

interaction between charged particles in chemical bonds

A

approaching nuclei and electrons attract each other

2 nuclei repel each other

2 electron clouds repel each other

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18
Q

bond length

bond distance

A

distance between 2 bonded atoms at their minimum potential energy

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19
Q

bond energy

A

energy (in kJ⁄mol) required to break a chemical bond and form neutral isolated atoms

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20
Q

octet

A

4 pairs of electrons

8 electrons

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21
Q

octet rule

A

states that chemical compounds tend to form so that each atom has an octet of electrons in its highest energy level by gaining, losing, or sharing electrons

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22
Q

exceptions to the octet rule

A

hydrogen and helium are stable with 2 electrons (1 pair of electrons)

aluminum and boron are stable with 6 electrons (3 pairs of electrons)

some elements can be surrounded with more than 8 electrons when combined with highly electronegative elements

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23
Q

electron-dot notation

A

electron-configuration notation in which only the valence electrons of an atom of a particular element are indicated by dots placed around the element’s symbol

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24
Q

explanation of a Lewis structure

A

atomic symbols represent nuclei and inner-shell electrons

dot-pairs or dashes between two atomic symbols represent electron pairs in covalent bonds

dots adjacent to only one atomic symbol represent unshared electrons

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25
unshared pair lone pair
pair of electrons that is not involved in bonding and that belongs exclusively to 1 atom
26
structural formula
indicates the kind, number, arrangement, and bonds of atoms in a molecule does not indicate the unshared pairs of atoms in a molecule
27
single bond single covalent bond
covalent bond produced by sharing 1 pair of electrons between 2 atoms represented by 2 dots in between 2 atoms' atomic symbols or by 1 dash
28
multiple bonds multiple covalent bonds
covalent bonds produced by sharing more than 1 pair of electrons between 2 atoms
29
types of multiple bonds
double bond triple bond
30
double bond double covalent bond
covalent bond produced by the sharing of 2 pairs of electrons between 2 atoms represented by 2 side-by-side pairs of dots or by 2 parallel dashes
31
triple bond triple covalent bond
covalent bond produced by the sharing of 3 pairs of electrons between 2 atoms represented by 3 side-by-side pairs of dots or by 3 parallel dashes
32
determining factor of need for multiple bonds
satisfaction of the octet rule for each atom
33
resonance
bonding in molecules or ions that cannot be correctly represented by a single Lewis structure
34
covalently bonded networks
compounds which can be pictured as continuous 3-dimensional networks of bonded atoms does not contain individual molecules
35
ionic compound
compound composed of positive and negative ions that are combined so that the numbers of positive and negative charges are equal
36
formula unit
simplest collection of atoms from which an ionic compound's formula can be established
37
crystal lattice
orderly arrangement of ions in an ionic compound which minimizes their potential energy maximizes attraction and minimizes repulsions of ions
38
lattice energy
energy released when 1 mole of an ionic compound is formed used to determine strength of ionic bonds
39
electrostatic forces
force that holds ions together in ionic compounds very strong overall attraction between positive and negative charges
40
strength of forces of covalent bonding (in comparison to forces of ionic bonding)
very weak
41
polyatomic ion
charged group of covalently bonded atoms
42
delocalized electrons sea of electrons
electrons that do not belong to any 1 atom move freely about metal's network of empty atomic orbitals
43
metallic bond
chemical bond that results from the attraction between metal atoms and the surrounding delocalized electrons does not result in the formation of a compound
44
metallic bond strength
varies directly with nuclear charge of metal atoms and number of electrons in the metal's electron sea directly proportional to heat of vaporization
45
metal's heat of vaporization
amount of heat required to vaporize a metal used to determine of strength of bonds
46
molecular polarity
uneven distribution of molecular charge determined by polarity of each bond with the geometry of the molecule
47
VSEPR
valence-shell electron pair repulsion
48
VSEPR theory
states that repulsion between the sets of valence-level electrons surrounding an atom causes these sets to be oriented as far apart as possible useful for explaining the structure of molecules helps to identify polarity of compounds
49
hybridization
combination of 2 or more orbitals of similar energy in same atom rearrangement of orbitals when an atom forms covalent bonds causes an overall reduction of energy
50
hybrid orbitals
orbitals of equal energy produced by the combination of two or more orbitals on the same atom
51
determining factor of number of hybrid orbitals
number of orbitals that have combined
52
intermolecular forces
forces of attraction between molecules
53
dipole
created by equal but opposite charges that are separated by a short distance represented by an arrow with a head pointing toward the negative pole and a crossed tail situated at the positive pole
54
direction of a dipole
from the dipole's positive pole to its negative pole
55
dipole-dipole force
intermolecular force between polar molecules due to attraction of opposite dipoles of each molecule lining up strong intermolecular force
56
hydrogen bonding
type of dipole-dipole intermolecular force between a hydrogen atom and a highly electronegative atom strongest intermolecular force results in very high melting point and boiling point usually represented by dotted lines connecting the hydrogen to the unshared electron pair of the electronegative atom
57
London dispersion force
intermolecular force between nonpolar molecules due to motion of electrons creating instantaneous poles on no polar molecules very weak intermolecular force results in very low melting point and boiling point causes compounds to be gaseous at room temperature
58
location of electrons in covalent bonds
overlapping orbitals
59
resonance structure
a separate (though equally valid) electron dot notation of a molecule or polyatomic ion
60
purpose of resonance structures
equalize bond lengths create stronger bonds
61
determining factor of need for resonance structures
presence of a double bond and a single bond on the same atom
62
coordinate bond dative covalent bond
covalently bonded pair of electrons which has been provided by only one of the bonding atoms
63
sp3 hybridization
hybridization of 1 s orbitals and 3 p orbitals forms no non-hybridized orbitals oriented in a tetrahedral geometry
64
sp2 hybridization
hybridization of 1 s orbital with 2 p orbitals forms 1 non-hybridized orbital oriented in a trigonal planar geometry
65
sp hybridization
hybridization of 1 s orbital and 1 p orbital forms 2 non-hybridized orbitals oriented in a linear geometry
66
linear geometry
molecular geometry of two or three atoms with bond angles of 180°
67
general formulas of linear geometry
AB[E3] AB2
68
bent geometry angular geometry
molecular geometry of three atoms with bond angles of 105°
69
general formulas of bent geometry general formulas of angular geometry
AB2[E] AB2[E2]
70
trigonal planar geometry
molecular geometry of four atoms with bond angles of 120°
71
general formula of trigonal planar geometry
AB3
72
tetrahedral geometry
molecular geometry of five atoms with bond angles of 109.5°
73
general formula of tetrahedral geometry
AB4
74
trigonal pyramidal geometry
molecular geometry of a five atoms with bond angles of 107°
75
general formula of trigonal pyramidal geometry
AB3[E]
76
trigonal bipyramidal geometry
molecular geometry of six atoms with bond angles of 120°
77
general formula of trigonal bipyramidal geometry
AB5
78
octahedral geometry
molecular geometry of seven atoms with bond angles of 90°
79
general formula of octahedral geometry
AB6
80
steps to determine molecular geometry for covalent compounds
draw electron dot diagram determine polarity of bonds based on electronegativity difference of bonded atoms count lone pairs on central atom determine whether or not compound has symmetry identify VSEPR shape determine molecular polarity
81
polarity of charged particles
no polarity
82
state of ionic compounds
solid at room temperature
83
melting point of ionic compounds
high
84
boiling point of ionic compounds
high
85
solubility in water of ionic compounds
soluble in water
86
conductivity of ionic compounds
conductive of electricity in liquid state (or when dissolved in water)
87
solubility in alcohol of ionic compounds
mostly insoluble in alchohol
88
state of polar covalent compounds
solid, liquid, or gas at room temperature
89
melting point of polar covalent compounds
low
90
boiling point of polar covalent compounds
low
91
solubility in water of polar covalent compounds
soluble in water
92
conductivity of polar covalent compounds
possibly conductive of electricity
93
solubility in alcohol of polar covalent compounds
possibly insoluble in alcohol
94
state of nonpolar covalent compounds
mostly liquid or gas at room temperature
95
melting point of nonpolar covalent compounds
low
96
boiling point of nonpolar covalent compounds
low
97
solubility in water of nonpolar covalent compounds
insoluble in water
98
conductivity of nonpolar covalent compounds
not conductive of electricity
99
solubility in alcohol of nonpolar covalent compounds
mostly soluble in alcohol

×Chemistry 1 (20 decks)

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