Transcript Chemical Bonds Chapter 8  A Chemical Bond is a link between atoms.  An Ionic Bond is the electrical attraction between.

Chemical Bonds
Chapter 8
 A Chemical Bond is a link between atoms.
 An Ionic Bond is the electrical attraction between the
opposite charges of cations and anions.
 A Lewis symbol consists of the chemical symbol of an
element and a dot for each of its valence electrons.Example
He:
 The formation of ionic bonds is represented in terms of
Lewis symbols by the loss or gain of electrons until both
species have reached an octet of electrons.
 The tendency to form cations two units lower in charge than
expected from the group number is called the inert pair
effect.Example consider the group 13 elements Al and
Indium. Al forms Al+3 but Indium forms In+1 and In+3. Group
14 has Pb that forms Pb+2 oxide when heated and tin forms
tin (IV)oxide when heated.
Lattice enthalpy
 A measure of the attraction between ions is the lattice
enthalpy, the enthalpy change per mole of formula units
when a solid is broken up into a gas of widely separated ions.
All lattice enthalpies are positive. Heat equal to the lattice
enthalpy is released when the solid lattice forms from
gaseous ions.
Lattice enthalpies and Born-Haber
Cycle
 The lattice enthalpy for a particular ionic compound is defined as ∆H
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for the process .
This cannot be measured directly as it is not possible to get this reaction
to happen on its own, without many other reactions happening around
it.
Separate ions cannot be brought together in this way. However, we can
use other pieces of experimental data to calculate this result. This is
known as a Born-Haber Cycle.
In a Born-Haber cycle, we imagine that we break apart the elements into
atoms, ionize the atoms, combine the gaseous ions to form the ionic
solid, then form the elements again from the ionic solid. Only the lattice
enthalpy, the enthalpy of the step in which the ionic solid is formed from
the gaseous ions, is unknown.
The overall energy change for a complete Born Haber cycle is 0.
 Start with the elements in the
proportions in which they
appear in the compound and
atomize them. Write the
corresponding enthalpies of
formation of the gas phase
atoms next to the upward
pointing arrows.
 Form gaseous cations from
the metal atoms. This step
requires the ionization energy
of the metal and the sum of
the first and higher ionization
energies. The arrow points
upwards.
 Form gaseous anions from the
nonmetal atoms. The enthalpy
change of this step is called the
electron gain enthalpy ∆Heg°. It
is the negative of the electron
affinity. If the electron affinity is
positive, the electron gain
enthalpy is negative and the
corresponding arrow points
downwards as the energy is
released. If the electron affinity
is negative then the electron
gain enthalpy is positive and the
arrow points upward.
 Let the gas of ions form a solid
compound. This step is the reverse of the
formation of ions from the solid so its
enthalpy change is the negative of the
lattice enthalpy, -∆Hl. Denote it by an
arrow pointing downward, since the
formation of the solid is exothermic. This
is the unknown value in the cycle.
 Complete the cycle with the arrow from
the compound to the element; the
enthalpy change in this step is the negative
of the enthalpy of the formation of the
compound from its elements, ∆Hf°. The
arrow points up if the ∆Hf° is negative,
down if it is positive.
 Finally calculate ∆Hl, from
the fact that the sum of all
the enthalpy changes for
the complete cycle is 0.
Covalent Bonds
 A covalent bond is a pair of electrons shared between two
atoms.
Octet rule and Lewis structure
 In covalent bonds, atoms share electrons to reach a noble gas
configuration. Lewis called this the octet rule.The valence of
an element is the number of covalent bonds an atom of the
element forms.Consider molecular hydrogen, H2. Each atom
completes its helium like duplet by sharing its electron with
the other:
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Class Practice
 Write the Lewis structure for the compound HBr and state
how many lone pairs each atom in the compound possesses?
Lewis structure for polyatomic species
 To write the Lewis structure for polyatomic species we count
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the valence electron from all the atoms in the molecule. For
example for methane there are 8 valence electrons.
The next step is to arrange the dots representing the
electrons so that the carbon atom has an octet and each
hydrogen atom has a duplet.
A single shared pair of electrons is called a single bond.
Atoms can share two or more electron pairs.
Two shared electron pairs form a double bond, and three
shared electron pairs form a triple bond.
Bonds
Home work
 Page 358
 8.48,8.52
 Write a Lewis structure for the amide ion NH2− .
Resonance
 In some Lewis structure the multiple bonds can be written in
several equivalent locations. Consider the nitrate ion,NO3−.
 The three Lewis structures shown differ only in the position
of the double bond. All the three structures are valid. The
bonds in the nitrate ion have a character intermediate
between a pure single bond and a pure double bond.We
present this as a blend of all three Lewis structure.
Formal charge
 The formal charge gives an indication of the extent to which
atoms have gained or lost electrons in the process of covalent
bond formation. Structures with lowest formal charges are
likely to have the lowest energy.
 Formal charge=number of valence electron in the free
atom−(number of electrons present as lone
pairs−½(number of electrons shared in bonds)
 =V−(L+½S)
Class practice
 Write three plausible structures with different atomic
arrangements for the cyanate ion, NCO− , and suggest which
one is likely to be the most plausible structure.
Radical and Biradical
 A radical is a species with an unpaired electron; a biradical
has two unpaired electron
Lewis acids and bases
 When a coordinate covalent bond forms, one species
provides a lone pair and the other species accepts it.The
species that provides the lone pair is called as Lewis base and
the species that accepts it is called as Lewis acid.
 To introduce this new class of reactions, lets investigate the
molecular structure of the colorless gas boron
trifluoride,BF3. The Lewis structure indicates that the boron
atom has an incomplete octet: its valence shell consists of
only six electrons. The molecule could complete its octet by
sharing more electrons with fluorine, but fluorine has such a
high ionization energy that this arrangement is not likely.
 This boron octet can be completed if another atom or ion
with a lone paired electrons forms a bond by providing the
needed pair of electrons. Example BF4−( tetrafluroborate
anion) forms when boron trifluoride is passed over a metal
fluoride. Now all the fluorine atoms have their normal
valence of 1and the boron atom has an octet.
Ionic versus Covalent bonds.
 Ionic and covalent are terms used to describe two extremes
of chemical bonds. When describing bonds with non metals
covalent bonds are good models and when a metal is involved
we say that it is an ionic bond. A covalent bond acquires some
ionic character if one atom has a greater electron
withdrawing power than the other atom. This electron
withdrawing power is called as electronegativity.When a
chemical bond forms between two atoms the atom with a
higher electronegativity pulls the atom with a lower
electronegativity.
 Electronegativity is a measure of the electron pulling power
of an atom on an electron pair in a molecule. Compounds
composed of elements with a large difference in
electronegativity (≥2) tend to have significant ionic character
in their bonding.
Class Practice
 In which of the following compounds do the bonds have
greater ionic character; NH₃ or NO₂?
 Indicate which atom in each compound has the partial
negative charge.
Home work
 Page359
 8.84,8.88
 Page 357
 8.36,8.38