Chapter 6 Fundamentals of Chemical Bonding 6.1 Overview of Bonding 6.2 Lewis Structures 6.3 Molecular Shapes: Tetrahedral Systems 6.4 Other Molecular Shapes

Chapter 6Fundamentals of Chemical Bonding

6.1 Overview of Bonding

6.2 Lewis Structures

6.3 Molecular Shapes: Tetrahedral Systems

6.4 Other Molecular Shapes

6.5 Properties of Covalent Bonds

Chemistry, 2nd Canadian Edition ©2013 John Wiley & Sons Canada, Ltd.


Learning objective:

Use the concept of electronegativity to determine the polarity of a chemical bond



Electrons and nuclei are continually moving.But they arrange themselves in ways that optimize the

net attractive forces among the electrons and the nuclei.

The net electrical energy can be calculated.


Schematic illustration of two electrons and two nuclei arranged so thatattractive coulombic interactions (blue lines) are greater than repulsivecoulombic interactions (red lines).


Chemical Bond Formation

Electrons 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.


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.


Example 6 -1

The bond length of molecular fluorine is 142 pm, and the bond energy is 155 kJ/mol. Draw a figure similar to Figure 6 – 2 that includes both F2 and H2. Write a caption for the figure that summarizes the comparison of these two diatomic molecules.


Let’s Look at F2

The fluorine atom has 7 valence electrons (2s22p5)By gaining an electron, it will become isoelectronic with

neon (2s22p6)If two fluorine atoms come together, they can share the

8th electron.


What happens to the orbitals with nonbonding electrons?

The orbitals are still there!The orbitals are still there!Chemistry, 2nd Canadian Edition ©2013 John Wiley & Sons Canada, Ltd.

Let’s Look at F2

Unequal Electron Sharing

A pure covalent bond occurs only when two identical atoms are bonded: N2, H2, F2, etc.

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.


The 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 we determine which atom has the partial negative charge and which atom has the partial positive charge?


Polar Bonds

Electronegativity is the Answer!

Electronegativity – the ability to attract bonding electrons.

Denoted by the Greek symbol chi, When two atoms have different electronegativities, the

bond between them is polar.The bigger the difference in electronegativities, the

more polar the bond.



Trends in Electronegativity

Electronegativity Increases

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.


Example 6 - 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.



Learning objective:

Draw optimized Lewis structures of covalent compounds, including resonance structures



Convenient representations of valence electrons Consists of the chemical symbol for the element plus a

dot for each valence electron.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


The Conventions

Follow the steps for drawing the Lewis Dot Structure of HF


The Conventions



The Conventions



The Conventions



The Bonding Framework

a. An outer atom bonds to only one other atom. An inner atom bonds to more than one other atom

b. Hydrogen atoms are always outer atoms.c. In inorganic compounds, outer atoms other than

hydrogen usually are the ones with the highest electronegativities.

d. The order in which atoms appear in the formula often indicates the bonding pattern

e. 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.


Building the Lewis Structure

1. Count the valence electrons.

2. Assemble the bonding framework, placing two electrons per bond.

3. Complete the octets on each outer atom, except H.

4. Assign the remaining electrons to inner atoms.

5. Optimize electron configurations of the inner atoms.

6. Identify equivalent or near-equivalent Lewis structures.


e.g. PCl3

Bonding Pairs

Lone Pairs (nonbonding electrons)

5 + (3 x 7) = 26 e-


Example 6 - 3

Determine the provisional Lewis structure of the BF4-

anion.


Example 6 - 4

Determine the provisional Lewis structure of diethylamine, (CH3CH2)2NH.


Optimizing the Structure

Step 5: Optimize 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.


Example 6 - 5

Aqueous solutions of formaldehyde, H2CO, are used to preserve biological specimens. Determine the Lewis structure of formaldehyde.


Example 6 - 6

Acrylonitrile, H2CCHCN, is used to manufacture polymers for synthetic fibers. Draw the Lewis structure of acrylonitrile.


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.


Example 6 - 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


Formal Charge

The difference between the number of valence electrons in the free atom and the number of electrons assigned to that atom in the Lewis structure.

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.

FC (Valence electrons in the free atom) -

(Valence electrons assigned to that atom in the Lewis structure)


Example 6 - 8

As described in Chapter 2, sulphur dioxide, a by-product of burning fossil fuels, is the primary contributor to acid rain. Determine the Lewis structure of SO2.


Example 6 - 9

Acetic Acid (CH3CO2H, a carboxylic acid) is an important industrial chemical and is the sour ingredient in vinegar. Build its Lewis structure.


Resonance Structures

Step 6: Identify equivalent or near-equivalent Lewis structures

Let’s look at nitrate, NO3-





Example 6 - 10

Determine the Lewis structure of dihydrogen phosphate, H2PO4

-.


Determine the Lewis structure of dinitrogen oxide (NNO), a gas used as an anaesthetic, a foaming agent, and a propellant for whipped cream.

Example 6 - 11


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.



Learning objective:

Recognize the importance of the tetrahedral shape in molecules


Molecules have three dimensional shapes.The 3-D shapes define the properties of the molecules.How do we predict the shapes?VSEPR Theory – valence shell electron-pair repulsion

theory Electron pairs in the outer shell of an atom repel one another

and end up as far away from each other as possible.



Let’s Take a Step Back…

Molecules have 3-D shapes because orbitals have 3-D shapes.

Let’s look at methane:


In the plane of the paper, it looks like the bond angles are 90°, but, we know that the molecule exists in three dimensions.

The shape is called tetrahedral and has bond angles of 109.5°.


Methane, CH4

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.


The VSEPR Model

First some definitions: Electron 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.


All molecules above have the samesteric number or electron group geometry

(3-D arrangement of the valence shell electron groups)Chemistry, 2nd Canadian Edition ©2013 John Wiley & Sons Canada, Ltd.

Molecular Shape

Molecular Shape

The molecular shape describes how the ligands (not the electron groups) are arranged in space.



1. Determine the Lewis structure.

2. Use the Lewis structure to find steric numbers for inner atoms.

3. Determine the electron group geometries from the steric

numbers.

4. Use the ligand count to derive molecular shapes from electron

group geometries.

Determining Molecular Shape

Example 6 - 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.


Example 6 - 13

Describe the shape of hydroxylamine, HONH2.


Silicon

Silicon displays tetrahedral shape to virtually all of its stable compounds.

95% of crustal rock and its various decomposition products are composed of silicon oxides.

The principle oxide of silicon is silica, with empirical formula SiO2.



Learning objective:

Use the VSEPR model to predict the shapes of molecules with steric numbers 2, 3, 5 and 6



Steric Number 2: Linear Electron Group Geometry

Steric Number 3: Trigonal Planar Electron Group Geometry


6.4 Other Molecular Shapes (cont.)

Steric Number 5: Trigonal Bipyramidal Electron Group Geometry


6.4 Other Molecular Shapes (cont.)

Steric Number 6: Octahedral Electron Group Geometry


Example 6 - 15

Describe the geometry and draw a ball-and-stick sketch of Xenon tetrafluoride.



Learning objective:

Understand the factors that influence bond angles, lengths and energies



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.


Example 6 - 16

Experiments show that sulphur tetrafluoride has bond angles of 86.9° and 101.5°.

Give an interpretation of these bond angles.


Dipole Moments

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,



Example 6 - 17

Does either ClF5 or XeF4 have a dipole moment?


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.


Table 6 – 1: Average Bond Lengths


Table 6 – 1: Average Bond Lengths


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.


Example 6 - 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.


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.


Table 6 – 2 Features of Molecular Geometries


Table 6 – 2 Features of Molecular Geometries



Chapter 6 Visual Summary









Documents

Chapter 6 Fundamentals of Chemical Bonding 6.1 Overview of Bonding 6.2 Lewis Structures 6.3 Molecular Shapes: Tetrahedral Systems 6.4 Other Molecular Shapes