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3.S: Chemical Bonding and Nomenclature (Summary)
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In
covalent bonding
, electrons are
shared
between atoms. In
ionic bonding
, electrons are transferred from one atom to another. Compounds that are formed using only covalent bonds are termed
molecular compounds
.
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The
outermost
electron level in any atom is referred to as the
valence shell
. The electron configuration of the valence shell of an atom can be shown graphically using a
Lewis diagram
(or
electron-dot structure
). The arrangement of “dots” around the chemical symbol for the element are shown singly up through four electrons, and then paired until eight electrons are present.
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To form ions from individual elements, electrons are added or subtracted from the valence shell in order to completely fill the shell with eight electrons (the
octet rule
). The charge on the ion reflects the electrons added or removed. Representative elements through Group
IIIA
will lose electrons to form cations, while those in groups
IVA
–
VIIA
will gain electrons and form anions.
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A
covalent bond
is constructed in a Lewis diagram by pairing a set of unpaired electrons from two different atoms. For the purposes of the “octet rule”, a pair of shared electrons is counted as two electrons for each atom. Multiple covalent bonds (
double bonds
and
triple bonds
) are used, if necessary, to give each bonded atom a full octet (except, of course, for helium and hydrogen). When two or more atoms are bonded together utilizing covalent bonds, the compound is referred to as a
molecule
.
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As a rule of thumb,
ionic
bonds will be formed whenever the compound contains a
metal
.
Covalent
bonding will be observed in compounds containing only
semimetals
or
nonmetals
.
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Groups of covalently bonded semimetals or nonmetals which are charged are called
polyatomic ions
. Common examples include sulfate dianion, nitrate anion, phosphate trianion, etc. These polyatomic ions are commonly paired with metals forming ionic compounds.
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Many (but not all) polyatomic ions can be drawn in two or more equivalent Lewis representations. These are called
resonance forms
of the ion. The actual electronic structure of the ion is a combination of these Lewis structures and is called the
resonance hybrid
.
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The
electronegativity
of an element is a measure of the tendency of that element to attract electrons towards itself. Electronegativities range from 0.6 to 4.0, with fluorine as the most electronegative element (a value of 4.0). The general trend in the periodic table is for electronegativity to increase from the lower left-hand corner (Fr) to the upper right-hand corner (F).
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Covalent bonds formed between atoms with different electronegativity will be
polarized
with the greatest electron density localized around the most electronegative atom. The effect of electronegativity on electron distribution within a molecule can be shown using a computer-calculated
electrostatic potential map
where colors are used to represent electron density.
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Elements in periods 3 – 7 can accommodate more that eight electrons in there valence shells. This phenomena is called
valence shell expansion
and molecules involving these elements may have 10 – 14 valence electrons in properly drawn Lewis diagrams. Exceptions to the “octet rule” also exist where the valence shell contains less than eight electrons, or contains unpaired electrons.
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When naming simple, binary
ionic
compounds, the cation is named first using the name of the element, followed by the anion, where the suffix
ide
is added to the root name of the element. Multipliers are
not used
. For transition metals in which the metal can assume a variety of oxidation states (different positive charges), the charge of the metal ion is shown in the name using Roman numerals, in parenthesis, following the name of the element (i.e., iron (III) chloride).
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When naming simple binary
molecular
compounds (compounds containing only covalent bonds) the
least electronegative
element is (generally) named first, followed by the second element, where the suffix
ide
is again added to the root name of the element. In molecular compounds multipliers are used to indicate the number of each atom present (mono-, di-, tri-, tetra-, etc.) with the exception that
mono
is not used for the first element in the compound.