VSEPR Theory. Each group of valence electrons around a central atom is located as far as possible from the others, to minimize repulsions.

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VSEPR Theory

 Electron pairs surrounding an atom repel each other. This is referred to as

Valence Shell Electron Pair Repulsion (VSEPR) theory.

 NOT based on how many electron pair there are, but where they are.  Where they are: electron groups (charge clouds)

 Each group of valence electrons around a central atom is located as far as possible from the others, to minimize repulsions.

C H H H H O C O C O N H H H O H H C N H

What counts as an electron group?

1) A connection between the atom of interest and any other atom a) single, double, and triple bonds each count one electron group

2) A lone pair on the atom of interest

S S O O O O H H

VSEPR Theory

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VSEPR Theory

 The electron pair geometry gives the arrangement of atoms AND the lone pair electrons around the central atom.

 There are five basic arrangements of electron groups around a central atom.

VSEPR Theory

 The electron pair geometry gives the arrangement of atoms AND the lone pair electrons around the central atom.

 There are five basic arrangements of electron groups around a central atom.

 The molecular shapeis the three–dimensional arrangement of nuclei joined by the bonding groups. This is defined only by the relative positions of the nuclei. 3

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For CO₂, the central C atom has two atoms attached (the two oxygen atoms) and has no lone pair.

 2 electron groups –  Hybridization –

 The electron pair geometry is –  The molecular shape is –

 Bond angle is –

Linear Molecules

O

C

O

Trigonal Planar

 For H₂CO, the central C atom has three atoms attached and has no lone pair.  3 electron groups

 Hybridization

 The electron pair geometry is  The molecular shape is

 Bond angle is

C O

H H

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Trigonal Planar

 In ozone, O₃, the central O atom has two atoms attached (the other 2 oxygens) and has one lone pair.

 3 electron groups  Hybridization –

 Bond angle is O O O S O O

Tetrahedral

 In SiCl₄, the central Si atom has four atoms attached (four chlorine atoms) and has no lone pair.

 4 electron groups  Hybridization

 The electron pair geometry is

Cl Si Cl Cl Cl 7

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Tetrahedral

 In NH₃, the central N atom has three atoms attached (three hydrogen atoms) and has one lone pair.

 Bond angle is

N

H

H H

Tetrahedral

 In SCl₂, the central S atom has two atoms attached (two chlorine atoms) and has two lone pair.

 Bond angle is

Cl S

Cl 9

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Trigonal Bipyramidal

 In PCl₅, the central P atom has five atoms attached (five chlorine atoms) and has no lone pair.

 The electron pair geometry is t

Cl P Cl Cl Cl Cl 11

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Trigonal Bipyramidal

 In SF₄, the central S atom has four atoms attached (four fluorine atoms) and has one lone pair.

 Bond angles F S F F F 13

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Trigonal Bipyramidal

 In BrF₃, the central Br atom has three atoms attached (three fluorine atoms) and has two lone pair.

 Bond angles Br F F F

Trigonal Bipyramidal

 In XeF₂, the central Xe atom has two atoms attached (two fluorine atoms) and has three lone pair.

 The electron pair geometry is

Xe F

F

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Octahedral

 In SF₆, the central S atom has six atoms attached (six fluorine atoms) and has no lone pair.

 Bond angles S F F F F F F

Octahedral

 In BrF₅, the central Br atom has five atoms attached (five fluorine atoms) and has one lone pair.

 Bond angles Br F F F F F 17

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Octahedral

 In XeF₄, the central Xe atom has four atoms attached (four fluorine atoms) and has two lone pair.

 Bond angles Xe F F F F 19

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Electronegativity and Polarity

A covalent bond in which the shared electron pair is not shared equally, but remains closer to one atom than the other, is a polar covalent bond.

If “X” and “Y” share bonding e-_equally:

If “X” and “Y” do NOT share bonding e-_equally:

Unequal sharing of bonding e-_{leads to polar covalent bonds}

Polar Covalent Bonds

Polar: having poles

N S

One end opposite from the other

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Electronegativity and Polarity

A covalent bond in which the shared electron pair is not shared equally, but remains closer to one atom than the other, is a polar covalent bond.

Unequal sharing of electrons causes the more electronegative atom of the bond to be partially negative and the less

electronegative atom to be partially positive.

The ability of an atom in a covalent bond to attract the BONDING electrons towards itself is called its electronegativity.

If electronegativity (EN) difference is:

Electronegativity difference is < 0.5 then the bond is considered to be pure covalent

C

S

Br Br

I

H

2.5 – 2.5 = 0 2.5 – 2.5 = 0 2.8 – 2.8 = 0 2.5 – 2.1 = 0.4

2.0 > electronegativity difference ≥ 0.5 then the bond is considered to be polar covalent

C O

F

H

N C

Si

O

Polarity Revisited

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The lowercase Greek letter delta, , is used to indicate a polar bond.

The MORE EN element has extra e–_{, so it is negative and is indicated by the}

symbol –_.

The LESS EN element is short of e–_{, so it is positive and is indicated by the}

symbol +_.

H–Cl

Electronegativity and Polarity

Give delta notation and polarity arrows for the following:

C

O

F

H

N

C

Si

O

Electronegativity and Polarity

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Molecular Polarity

A molecule is polar if

- it contains one or more polar bonds and

- the individual bond dipoles do not cancel (shape).

Overall molecular polarity depends on both shape and bond polarity.

The polarity of a molecule is measured by its dipole moment (μ), which is given in the unit debye (D).