CHAPTER 8 – Bonding: General Concepts
• Molecular structure is an important key to understanding properties
of matter.
Molecular structure deals with:
1)
Shape of molecules
2)
Distribution of charge
Chemical Bond -
Three Extreme types
1.) Ionic Bonds
2.) Covalent Bonds
3.) Metallic Bonds
Two main reasons why Chemical events occur:
1.) Movement toward ______________________
2.) Movement toward ______________________
• THUS the main reason why chemical bonds form is because the system is at
__________________ after the bond forms.
Electronegativity:
Trends:
Increases
Decreases
Electronegativities are assigned based on the comparison of expected bond energies to
actual bond energies.
Expected H-X bond energy = H-H bond energy + X-X bond energy
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Δ = (H-X)act – (H-X)exp
Relative electronegativities of H and X can be assigned from Δ values.
Ex. Order the following bonds according to polarity:
H-H, O-H, Cl-H, S-H, and F-H
1
100%
Covalent
100%
ionic
Ions: Electron Configuration and Sizes
• When two nonmetals react to form a ______________ bond, they ___________ electrons in a way that
completes the valence electron configurations of both atoms.
• When a nonmetal and a representative-group metal react to form a binary ___________ compound, the
ions form so that the valence electron configuration of the non metal achieve the electron configuration
of the next ______________ atom and the valence orbitals of the metal are __________________.
• solid ionic compounds contain a large collection of positive and negative ions packed together in a way
that minimizes the (-).. (-) and (+).. (+) repulsions and maximizes the (+)..(-) attractions.
Some exceptions: Sn2+/Sn4+
Pb2+/Pb4+
Tl+/Tl3+
Bi3+/Bi5+
Ionic Radii:
• Positive ions are _________than the metal atoms from which they are formed.
• Negative ions are _________than the nonmetal atoms from which they were formed.
• The greater the change in charge the greater the difference in size, thus Co3+ is ____________ than Co2+
isoelectronic ions – ions containing the same ___________________
Ex. Rank these in order of decreasing size: Se2-, Br-, Rb+, Sr2+
Ex. Choose the largest ion in each of the following groups:
a. Li+, Na+, K+, Rb+, Cs+
b. Ba2+, Cs+, I-, Te2In general, for a series of isoelectronic ions, the size ___________ as the nuclear charge Z increases
Formation of Binary Ionic Compounds:
Lattice energy –
M+(g) + X - (g)  MX (s)
• The energy associated with the formation of a solid ionic compound can be calculated by
breaking the reaction into steps (since energy is a state function) and then be summed.
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Note: The energy released when an electron is added to a fluorine atom to form the F- ion (328 kJ/mol) is
not enough to remove an electron from lithium (520 kJ/mol). That is, when a metallic atom reacts
with a nonmetallic atom to form separated ions.
• The main driving force for the formation of an ionic compound rather than a covalent
compound results from the strong mutual attractions among the positive and negative
ions in the solid. The lattice energy term is the dominant term.
Lattice Energy = k (Q1 Q2)
r
According to Coulomb’s Law, Lattice Energies Depend on:
1) the charges of the ions involved (Q1 and Q2)
2) the distance between the two nuclei (r)
a.k.a. the size of the ions
Identify which compound would have the larger lattice energy in each of the following:
1. NaF and KF
2. BeO and LiF
3. MgCl2 and BeCl2
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Partial Ionic Character of Covalent Bonds
% ionic character of the bond in diatomic molecule XY:
% ionic character =
measured dipole moment of X-Y x 100 %
calculated dipole moment of X+Y-
- Generally compounds are considered ionic if they have at least 50% ionic character.
Problem: This is measured from discrete molecules in the gaseous state, not solids.
- The book will consider anything ionic that conducts an electric current when melted.
• Ionic character __________________ with difference in electronegativity
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The Covalent Chemical Bond: A Model
Chemical bond - forces that cause a group of atoms to behave as a unit
• Bonds result from the tendency of a system to seek its
________________ possible energy
OR
• Bonds occur when collections of atoms are more stable (lower in
energy) than the separate atoms.
(A bond represents a quantity of energy obtained from the overall molecular energy of stabilization.)
• The model of the chemical bond has helped chemists to systematize the reactions of the
millions of existing compounds, but remember that the chemical bond is only a model.
Covalent Bond Energies and Chemical Reactions:
Bond Energies:
• The Bond energy is defined as H when one mole of bonds is broken in the
gaseous state.
• Energy is _____________ when a chemical bond is broken.
H =
EXO/ ENDO
• Energy is _____________ when a chemical bond is formed.
H =
EXO/ ENDO
An estimated H for a reaction can be found.
From Hrxn =
 B.E. bonds broken - B.E. bonds formed
Example: Estimate H for a reaction to form ammonia from its elements.
• The values in table 8.4 are averages thus when used to predict H’s of
reactions the answers are approximate.
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Ex.
CH3-CH2-CH2-O-H
O
||
vs. CH3-CH2- C-O-H
• Polar molecules have _______ bonds than non-polar molecules.
• Double and triple bonds are _______ than single bonds.
Bond Length: The distance between the centers of 2 bonded atoms
• Non-polar molecules:
• Polar molecules:
• Bond distance for double and triple bonds is ________ than single bonds.
The Localized Electron Bonding Model:
assumes that a molecule is composed of atoms that are
bound together by sharing pairs of electrons using the
____________________ of the bound atoms.
• electron pairs in the molecules are assumed to be localized on a particular
atoms or in the space between two atoms.
LE model has three parts:
1) Description of the valence electron arrangement in the molecule using Lewis
structures
2) Prediction of the geometry of the molecules using the valence shell electron
pair repulsion (VSEPR) model.
3) Description of the type of atomic orbitals used by the atoms to share electrons
or hold lone pairs.
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Rules for Lewis Dot Structures
1.) Count the number of valence electrons.
a. Valence e- equal group number of element.
b. Adjust for ion charges.
Negative Positive -
2.) Draw a skeleton structure for the species, joining atoms with single bond.
NOTE:
a. central atoms written first
If a molecule/ion
b. Peripheral atoms usually hydrogen, halogens and oxygen
consists of 3 atoms
connected in a line,
c. Isomers are possible C2H6O
they will be written in
that order. (As
opposed to listing the
central atom first.
Ex. SCN- CNO-,
NCO-, NSF
3.) Deduct 2 valence electrons for each single bond used in step 2.
a. Distribute the remaining electrons as unshared pairs so as to give
each atoms 8 e- if possible
b. Not enough electrons use multiple bonds -double & triple - (C, N, O, S)
CH4O
OCl-
NH4+
NO3-
CN-
• Multiple bonds are shorter than single bonds.
Resonance Structures:
• Resonance forms are used when observed bond lengths do not correlate with
Lewis dot structure.
1.) Resonance forms are obtained by moving electrons, not atoms.
2.) Resonance can be expected when it is possible to draw more than one
structure that follows the octet rules.
* The TRUE structure is the average “hybrid” of the various resonance forms.
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Exceptions to the Octet Rule
A.) Odd numbers of valence electrons
NO
NO2
B.) Group II and group III elements often have less than a full octet.
C.) Expanded Octet - Some elements when present as a central atom use
empty d orbitals along with the commonly used s and p orbitals. This
increases their capacity from 8 --> 10 or 12 e-.
• Elements that can do this are in:
rows 3, 4, & 5
& groups 5, 6, 7, & 8
• When writing Lewis structures, do not worry about which electrons come from which atoms in
a molecule. The best way to look at a molecule is to regard it as a new entity that uses all the
available valence electrons of the atoms to achieve the lowest possible energy.
• In a sense this approach corrects for the fact that the localized electron model
overemphasizes that a molecule is simply a sum of its parts, that is, that the atoms retain their
individual identities in the molecule.
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Molecular Structure: The VSEPR Model
VSEPR Model - Valence Shell Electron Pair Repulsion
Bond Angles and Shapes Electron groups stay away from each other as much as possible.
An Electron group may be:
1) a non-bonded pair
2) A bonded pair (single bonds)
3) A multiple bond (double and triple bonds)
Count electron groups about central atom
If 2 groups - it is linear (180°)
If 3 groups - bond angle = 120°
If all are bonding, then planar triangular
If 2 bonding and 1 non-bonding group
then bent @ = 120°
If 4 groups - bond angle = 109.5 °
all 4 bonding
3 bonding + 1 non-bonding
tetrahedral
pyramidal or tripod
2 bonding + 2 non bonding
bent @
Note: lone pairs require more room than bonding pairs and tend to compress the
angles between the bonding pairs.
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Covalent Bond Properties
Bond Polarity:
• Pure covalent bonds result from ________ sharing of electrons
• Polar Covalent Bonds result from ____________ sharing of electrons.
• Ionic Bonds result from __________________________
A molecule is polar if:
1. It has ________ bonds.
2. It has a geometry such that the polarity ____________________.
Formal Charge:
• When several nonequivalent Lewis structures can be drawn for a molecule,
formal charge is often used to choose the most appropriate structures(s).
Formal charge =
• The sum of the formal charges of all atoms in a given molecule or ion must
equal the overall charge on that species.
Calculate the formal charges for each atom in the following molecules/ions:
NH4+
SO2
PCl5
Calculate the formal charges for each atom in the two structures of SO42-.
• If nonequivalent Lewis structures exist for a species, those with formal charges closest to zero
and with any negative formal charges on the most electronegative atoms are considered to
best describe the bonding in the molecule or ion.
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