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CHAPTER OUTLINE���

  Electronegativity   Polarity & Electronegativity   Lewis Structures   Molecular Shapes

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ELECTRONEGATIVITY���

  Electronegativity (E.N.) is the ability of an atom involved in a covalent bond to attract the bonding electrons to itself.

  Linus Pauling derived a relative Electronegativity Scale based on Bond Energies.

Cs 0.7

F 4.0

Least electronegative

Most electronegative

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ELECTRONEGATIVITY���

Electronegativity increases

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BOND POLARITY &���ELECTRONEGATIVITY���

The more polar the

bond formed

Polarity is a measure of the inequality in the sharing of bonding electrons

The more different the

electronegativity of the elements

forming the bond

The larger the electronegativity

difference (ΔEN)

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POLARITY &���ELECTRONEGATIVITY���

As difference in electronegativity

increases

Bond polarity increases

Most polar

Least polar

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POLARITY &���ELECTRONEGATIVITY���

Electronegativity difference

Bond Type

ΔEN = 0 Non-polar covalent

0 < ΔEN <1.7 Polar covalent

1.7 < ΔEN Ionic

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H H

Hydrogen Molecule

The molecule is nonpolar covalent

Electronegativity 2.20

Electronegativity 2.20

POLARITY &���ELECTRONEGATIVITY���

ΔEN = 0

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H Cl

Hydrogen Chloride Molecule

Electronegativity 2.20

Electronegativity 3.16

The molecule is polar covalent

δ+ δ-

ΔEN = 0.96

POLARITY &���ELECTRONEGATIVITY���

9 Sodium Chloride

Na+ Cl-

Electronegativity 0.93

Electronegativity 3.16

The bond is ionic No molecule exists

ΔEN = 2.23

POLARITY &���ELECTRONEGATIVITY���

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SUMMARY���OF BONDING���

Ionic Bond (large ΔEN)

Covalent Bond (small to moderate ΔEN)

Non-polar (similar electronegativities)

Polar (moderate ΔEN)

EN > 1.7

EN = 0

0 < EN < 1.7

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COMPARING PROPERTIES���OF IONIC & COVALENT���

COMPOUNDS���

Ionic Covalent

Structural Unit Ions Molecules

Melting Point High Low

Boiling Point High Low

Solubility in H2O High Low or None

Electrical Cond. High None Examples NaCl, AgBr H2, H2O

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LEWIS ���STRUCTURES���

  Lewis structures use Lewis symbols to show valence electrons in molecules and ions of compounds.

  Lewis symbols for the first 3 periods of representative elements are shown below:

  In Lewis symbols, valence electrons for each element are shown as a dot.

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LEWIS ���STRUCTURES���

  In a Lewis structure, a shared electron pair is indicated by two dots between the atoms, or by a dash connecting them.

  Unshared pairs of valence electrons (called lone pairs) are shown as belonging to individual atoms or ions.

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LEWIS ���STRUCTURES���

  Writing correct Lewis structures for covalent compounds requires an understanding of the number of bonds normally formed by common nonmetals.

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LEWIS ���STRUCTURES���

  Covalent molecules are best represented with electron-dot or Lewis structures.

  Structures must satisfy octet rule (8 electrons around each atom).

  Hydrogen is one of the few exceptions and forms a doublet (2 electrons).

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LEWIS���STRUCTURES���

  Bonding electrons can be displayed by a dashed line.

  Non-bonding electrons must be displayed as dots.

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LEWIS���STRUCTURES���

  More complex Lewis structures can be drawn by following a stepwise method:

1. Count the number of electrons in the structure.

2. Draw a skeleton structure.

3. Connect atoms by bonds (dashes or dots).

4. Distribute electrons to achieve Octet rule.

5. Form multiple bonds if necessary.

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Example 1:���

Write Lewis structure for H2O

H2O = 8 electrons 2 (1) + 6 = 8 Step 1:

Step 2: H O H

Skeleton structure should be

symmetrical

Step 3:

4 electrons used 4 electrons remaining Step 4:

• •

• •

Octet rule is satisfied

Hydrogen has doublet

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Example 2:���

Write Lewis structure for CO2

CO2 = 16 electrons

4 + 2(6) = 16 Step 1:

Step 2:

O C O Skeleton structure should be

symmetrical

Step 3:

4 electrons used 12 electrons remaining

Step 4:

• •

• •

Octet rule is satisfied

• •

10 electrons used 6 electrons remaining

• •

• •

• •

Octet rule is NOT satisfied

Step 5:

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Example 3:���

Write Lewis structure for CO32-

CO32- = 24

electrons 4+3(6)+2 = 24 Step 1:

Step 2:

O C O

O Step 3:

Step 4:

• •

• •

• •

• •

• •

• •

• •

• •

• •

18 electrons remaining

12 electrons remaining 6 electrons remaining 0 electrons remaining

Octet rule is satisfied Octet rule is NOT satisfied

Step 5:

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Example 4:���

Write Lewis structure for NH3

NH3 = 8 electrons 5 + 3(1) = 8 Step 1:

Step 2:

H N H

H Step 3:

Step 4:

• •

Octet rule is satisfied

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Example 5:���

Write Lewis structure for ClO3-

ClO3- = 26

electrons 7+3(6)+1 = 26 Step 1:

Step 2:

O Cl O

O Step 3:

Step 4:

• •

• •

• •

• •

• •

• •

• •

• •

• •

20 electrons remaining

14 electrons remaining 8 electrons remaining 2 electrons remaining

Octet rule is satisfied

• • 0 electrons remaining

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EVALUATING LEWIS���STRUCTURES���

  When evaluating Lewis structures for correctness, two points must be considered:

1.  Are the correct number of electrons present in the structure?

2.  Is octet rule satisfied for all elements? (Hydrogen is an exception)

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Example 1:���

Determine if each of the following Lewis structures are correct or incorrect. If incorrect, rewrite the correct structure.

2(1) + 4 + 6 = 12

Octet is complete Doublets are complete

Octet is incomplete

Structure is incorrect

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Example 2:���

Determine if each of the following Lewis structures are correct or incorrect. If incorrect, rewrite the correct structure.

2(5) + 4(1) = 14

Structure is incorrect

Only 12 electrons shown

2 2

2 2 4

Structure has 14 electrons

Octets are complete

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MOLECULAR���SHAPES���

  The three-dimensional shape of the molecules is an important feature in understanding their properties and interactions.

  All binary molecules have a linear shape since they only contain two atoms.

  More complex molecules can have various shapes (linear, bent, etc.) and need to be predicted based on their Lewis structures.

  A very simple model , VSEPR (Valence Shell Electron Pair Repulsion) Theory, has been developed by chemists to predict the shape of large molecules based on their Lewis structures.

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MOLECULAR���SHAPES���

  Based on VSEPR, the electron pair groups in a molecule will repel one another and seek to minimize their repulsion by arranging themselves around the central atom as far apart as possible.

  Electron pair groups can be defined as any one of the following: bonding pairs non-bonding pairs

multiple bonds

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SUMMARY OF���VSEPR SHAPES���

Number of electron pair groups around

central atom Molecular

Shape Bond Angle

Examples

Bonding Non-bonding

2 0 Linear 180 CO2

3 0 Trigonal planar 120 BF3

2 1 Bent 120 SO2

4 0 Tetrahedral 109.5 CH4

3 1 Pyramidal 109.5 NH3

2 2 Bent 109.5 H2O