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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


• 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.

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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.


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.


Chemical Bonding








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


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.


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.


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.


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


Resonance Structures

Step 6 Identify equivalent or near-equivalent Lewis structures

Lets look at nitrate, NO3-


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.


9.3: 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


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°.


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 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.


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


Molecular Shape

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



Chemical Bonding








Steric Number 2: Linear Electron Group Geometry

Steric Number 3: Trigonal Planar Electron Group Geometry

Steric Number 5: Trigonal Bipyramidal Electron Group Geometry


Steric Number 6: Octahedral Electron Group Geometry



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,


Hydrogen Halides


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 9 – 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.


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 9 – 3 Features of Molecular Geometries


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.

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Documents

courtesy james gimzewski

filled 2p orbital

partial negative charge

partial positive charge

zurich research laboratory

provisional lewis structure

separation distance

bond energy