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Chapter 9Fundamentals of
Chemical Bonding
9.1 Overview of Bonding
9.2 Lewis Structures
9.3 Molecular Shapes: Tetrahedral Systems
9.4 Other Molecular Shapes
9.5 Properties of Covalent Bonds
Courtesy James Gimzewski, IBM Research, Zurich Research Laboratory
9.1 Overview of Bonding
• Electrons and nuclei are continually moving.• But, in the motion, they arrange themselves in ways that optimize
the net attractive forces among the electrons and the nuclei.• The net electrical energy can be calculated.
rqqkE lectrical
21e
Chemical Bond FormationElectrons and nuclei in a molecule balance all interactions to give the
molecule stability.
Balance is achieved when the electrons are concentrated between the nuclei.
The electrons are shared between the nuclei and this sharing is called a covalent bond.
9.1 Overview of Bonding
Covalent bond:The attractive force that holds together two atoms that
share one or more electron pairs.
Ionic bond:Each ion has its own noble gas electron configuration
Covalent bond:Each atom has a noble gas electron configuration
but shares electron pair(s) to do so
H. + .H H : H or H — H
The two dots or the straight line drawn between the two atoms represent the covalent bond that holds the atoms
together.
Every covalent bond has a characteristic
length that leads to maximum stability.
This is the bond length.
The electron cloud or charge density formed by the two electrons is concentrated in the region
between the two nuclei.
When bonding electrons are between two nuclei, both nuclei are attracted to the electrons
The electrons link the nuclei together; the electrons are the “glue” that bonds the atoms to
each other.
The Hydrogen Molecule
Bond Length and Bond Energy
Bond length – the separation distance where the molecule is most stable
Bond energy – the amount of stability at this separation distance, also known as the strength of the bond.
F2
The fluorine atom has 7 valence electrons 1s22s22p5
It wants to have 8 valence electrons
This is also the electron configuration of Ne
F F-
If two fluorine atoms come together, they can share the 8th electron.
How do they come together?
+
A half-filled2p orbitalfrom one F
atom overlapsa half-filled2p orbital
from the otherF atom
Electron clouds of individual atoms overlap to form covalent bonds.
What happens to the orbitals with nonbonding electrons?
The orbitals are still there!The orbitals are still there!
Unequal Electron SharingA pure covalent bond occurs only when two identical atoms are
bonded: N2
When two dissimilar atoms form a covalent bond, the electron pair is unequally shared, the bond is called a polar covalent bond
Therefore, the electrons are nearer to one of the atoms, and that atom acquires a partial negative charge.
And consequently the other atom has a partial positive charge.
Bond is referred to as polar and the molecule can be called a dipole (having two poles)
The Greek symbol delta “” is used to indicate partial charge
How do you determine which atom has the partial negative charge and which atom has the partial positive charge?
Electronegativity – the ability to attract bonding electrons.The bigger the difference in electronegativities
between two atoms, the more polar the bond.
Greek symbol chi,
Polar Bonds
Nonmetals are more electronegative than metals.
In general: the further apart the atoms are on the periodic table, the larger the difference in electronegativity.
And, the larger the difference in electronegativity, the more polar the bond.
Chapter 9Fundamentals of
Chemical Bonding
9.1 Overview of Bonding
9.2 Lewis Structures
9.3 Molecular Shapes: Tetrahedral Systems
9.4 Other Molecular Shapes
9.5 Properties of Covalent Bonds
Courtesy James Gimzewski, IBM Research, Zurich Research Laboratory
9.2 Lewis Structures
Convenient representations of valence electrons
Consists of the chemical symbol for the element plus a dot for each valence electron.
Element symbol represents the nucleus.
In normal circumstances, 2 electrons per side, 4 sides.
If all sides are full, 8 electrons are in the valence shell…this is called an octet
F1s22s22p5
+ e– F–
1s22s22p6
F2 versus N2
9.2 Lewis Structures
Follow the steps for drawing the Lewis Dot Structure of HF
The Bonding Framework
An outer atom bonds to only one other atom. An inner atom bonds to more than one other atom
Hydrogen atoms are always outer atoms.
In inorganic compounds, outer atoms other than hydrogen usually are the ones with the highest electronegativities.
The order in which atoms appear in the formula often indicates the bonding pattern
The hydrogen atoms appear first in the formula of oxoacid. Nevertheless, in almost all cases these acidic hydrogen atoms bond to oxygen atoms, not to the central atom.
9.2 Lewis Structures
Building the Lewis Structure
An outer atom other than hydrogen is most stable when it is associated with an octet of electrons 9.2 Lewis Structures
Optimizing the Structure
Step 5Optimize electron configurations of inner atoms.
Check to see if any inner atoms lacks an octet. If needed, move electrons from adjacent outer atoms to make double or triple bonds until the
octet is complete.
9.2 Lewis Structures
Beyond the Octet
Elements in the 3rd period or higher can have more than an octet if needed.
Atoms of these elements have valence d orbitals, which allow them to accommodate more than
eight electrons.
9.2 Lewis Structures
Formal ChargeThe difference between the number of valence electrons in the free
atom and the number of electrons assigned to that atom in the Lewis structure.
FC = formal charge; G.N. = Group Number#BE = bonding electrons; #LPE = lone pair electrons
If Step 4 leads to a positive formal charge on an inner atom beyond the second row, shift electrons to make double or triple bonds to
minimize formal charge, even if this gives an inner atom with more than an octet of electrons.
LPEBE ##21- G.N. FC
9.2 Lewis Structures
Resonance Structures
Step 6 Identify equivalent or near-equivalent Lewis structures
Lets look at nitrate, NO3-
9.2 Lewis Structures
9.2 Lewis Structures
Hints on Lewis Dot Structures
1. Octet rule is the most useful guideline.2. Carbon forms 4 bonds.3. Hydrogen typically forms one bond to other
atoms.4. When multiple bonds are forming, they are
usually between C, N, O or S.5. Nonmetals can form single, double, and triple
bonds, but not quadruple bonds.6. Always account for single bonds and lone
pairs before forming multiple bonds.7. Look for resonance structures.
9.2 Lewis Structures
9.3: Molecular Shapes: Tetrahedral Systems
Molecules have three dimensional shape.
The 3-D shape defines the properties of the molecules.
How do we predict the shape?
VSEPR Theory – valence shell electron-pair repulsion theory
Electron pairs in the outer shell of an atom try to get as far away from each other as possible
Why? Because like charges repel…they want to be far away from each other.
Let’s take a step back…
Molecules have 3-D shape
Why? Because the electrons orbit the nucleus in 3-D orbitals.
Let’s look at methane
9.3: Molecular Shapes: Tetrahedral Systems
In the plane of the paper, it looks like the bond angles are 90°.
But, we know that the molecule exists in three dimensions.
So, the bonds are really optimized around the central carbon.
The shape is called tetrahedral and has bond angles of 109.5°.
9.3: Molecular Shapes: Tetrahedral Systems
Carbon and the Tetrahedron
Hydrocarbons – molecules that contain only carbon and hydrogen
Alkanes – hydrocarbons in which each carbon atom forms bonds to four other atoms.
9.3: Molecular Shapes: Tetrahedral Systems
The VSEPR ModelElectron group – a set of electrons that occupies a particular region
around an atom.Ligand – an atom or a group of atoms bonded to an inner atom.Steric number – the sum of the number of ligands plus the number of
lone pairs; in other words, the total number of groups associated with that atom.
9.3: Molecular Shapes: Tetrahedral Systems
All molecules above have the same steric number.
Electron group geometry – is the 3-D arrangement of the valence shell electron groups, corresponds to the steric
number
9.3: Molecular Shapes: Tetrahedral Systems
Molecular Shape
The molecular shape describes how the ligands (not the electron groups) are arranged in space.
9.3: Molecular Shapes: Tetrahedral Systems
Chapter 9Fundamentals of
Chemical Bonding
9.1 Overview of Bonding
9.2 Lewis Structures
9.3 Molecular Shapes: Tetrahedral Systems
9.4 Other Molecular Shapes
9.5 Properties of Covalent Bonds
Courtesy James Gimzewski, IBM Research, Zurich Research Laboratory
9.4 Other Molecular Shapes
9.4 Other Molecular Shapes
Steric Number 2: Linear Electron Group Geometry
Steric Number 3: Trigonal Planar Electron Group Geometry
Steric Number 5: Trigonal Bipyramidal Electron Group Geometry
9.4 Other Molecular Shapes
Steric Number 6: Octahedral Electron Group Geometry
9.4 Other Molecular Shapes
9.5 Properties of Covalent Bonds
Bond angles – each of the steric groups results in well-defined bond angles.
When the steric number of an atom changes, bond angles change exactly as the model predicts.
Lone pairs in a molecule cause bond angles to be a few degrees smaller than predicted for symmetrical geometry.
Dipole moments
Polar bonds
Polar bonds can result in polar molecules, depending on the molecule’s geometry
A polar molecule will align itself in an electric field
The extent to which the molecules align in a field is referred to as the dipole moment and has the Greek symbol mu,
9.5 Properties of Covalent Bonds
Hydrogen Halides
9.5 Properties of Covalent Bonds
Bond Lengths and Energies
Two important properties of bonds to study:
Bond length – the nuclear separation distance where the molecule is most stable.
Bond energy – the stability of a chemical bond
9.5 Properties of Covalent Bonds
Table 9 – 1: Average Bond Lengths
9.5 Properties of Covalent Bonds
What affects bond length?
1. The smaller the principle quantum numbers of the valence orbitals, the shorter the bond.
2. The higher the bond multiplicity, the shorter the bond.
3. The higher the effective nuclear charge of the bonded atoms, the shorter the bond.
4. The larger the electronegativity difference, the shorter the bond.
9.5 Properties of Covalent Bonds
Bond Energy
1. Bond strength increases as more electrons are shared between the atoms
2. Bond strength increases as the electronegativity difference (∆χ) between bonded atoms increases.
3. Bond strength decreases as bonds become longer.
9.5 Properties of Covalent Bonds
Table 9 – 3 Features of Molecular Geometries
9.5 Properties of Covalent Bonds
Example 9 -1
The bond length of molecular fluorine is 142 pm, and the bond energy is 155 kJ/mol. Draw a figure similar to Figure 9 – 2 that includes both F2 and H2. Write a caption for the figure that summarizes the comparison of these two diatomic molecules.
Example 9 - 2
Use the periodic table, without looking up electronegativity values, to rank each set of three bonds from least polar to most polar: (a) S – Cl, Te – Cl, Se – Cl; and (b) C – S, C – O, and C – F.
Example 9 - 3
Determine the provisional Lewis structure of the BF4
- anion.
Determine the provisional Lewis structure of diethylamine, (CH3CH2)2NH.
Example 9 - 4
Example 9 - 5
Aqueous solutions of formaldehyde, H2CO, are used to preserve biological specimens. Determine the Lewis structure of formaldehyde.
Example 9 - 6
Acrylonitrile, H2CCHCN, is used to manufacture polymers for synthetic fibers. Draw the Lewis structure of acrylonitrile.
Example 9 - 7
Chlorine trifluoride is used to recover uranium from nuclear fuel rods in a high temperature reaction that produces gaseous uranium hexafluoride
2 ClF3 (g) + U (s) UF6 (g) + Cl2 (g)
Determine the Lewis structure of ClF3
Example 9-8
As described in Chapter 5, sulfur dioxide, a by-product of burning fossil fuels, is the primary contributor to acid rain. Determine the Lewis structure of SO2.
Example 9 - 9
Acetic Acid (CH3CO2H, a carboxylic acid) is an important industrial chemical and is the sour ingredient in vinegar. Build its Lewis structure.
Example 9 - 10
Determine the Lewis structure of dihydrogen phosphate, H2PO4
-.
Determine the Lewis structure of dinitrogen oxide (NNO), a gas used as an anesthetic, a foaming agent, and a propellant for whipped cream.
Example 9 - 11
Example 9 - 12
Describe the shape of the hydronium ion (H3O+). Make a sketch of the ion that shows the three-dimensional shape, including any lone pairs that may be present.
Describe the shape of hydroxylamine, HONH2.
Example 9 - 13
Example 9 - 15
Describe the geometry and draw a ball-and-stick sketch of Xenon tetrafluoride.
The third molecular shape arising from an octahedron is exemplified by chlorine pentafluoride, ClF5. Describe the geometry and draw the ball-and-stick of chlorine pentaflouride.
Extra Practice Exercise 9 - 15
Example 9 - 16
Experiments show that sulfur tetrafluoride has bond angles of 86.9° and 101.5 °. Give an interpretation of these bond angles
Example 9 - 17
Does either ClF5 or XeF4 have a dipole moment?
Use molecular symmetry to determine if ethane (C2H6) and ethanol (C2H5OH) have dipole moments.
Extra Practice Exercise 9 - 17
Comparison of bonding between two carbon and two silicon atoms.
Example 9 - 18
What factors account for each of the following differences in bond length?
a. I2 has a longer bond than Br2.b. C – N bonds are shorter than C – C bonds.c. H – C bonds are shorter than C ≡ Od. The carbon – oxygen bond in
formaldehyde, H2C=O, is longer than the bond in carbon monoxide, C ≡ O.