1 Chapter 8: Ionic and Covalent Bonding RVCC Fall 2009 CHEM 103 – General Chemistry I Chemistry: The Molecular Science, 3 rd Ed. by Moore, Stanitski, and.

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Chapter 8:

Ionic and Covalent Bonding

RVCC Fall 2009CHEM 103 – General Chemistry I

Chemistry: The Molecular Science, 3rd Ed. by Moore, Stanitski, and Jurs

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Bonding – What holds atoms together?

Octet rule: Octet rule: To form bonds, atoms gain, lose, or share e- to achieve a valence shell of 8 (or isoelectronic with a noble gas).

Ionic bond – an electrostatic attraction between a cation and an anion that forms when electrons transfer from one atom to another.

Covalent (Molecular) bond – the net attractive force that results from the sharing of electrons between atoms.

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Ionic Bonds

An ionic bond is formed by the transfer of electrons from one atom (metal with low EA) to another (nonmetal with high EA). The resultant ions are held together by electrostatic attraction.

Na.:

Cl: . :

Na+ Cl-

[Ne]3s1 [Ne]3s23p5

[Ne] [Ne]3s23p6 = [Ar]

Each atom has satisfied the octet rule.

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Ionic Bonds

::

F.:

:

:F .: Mg. .

Mg F

:

:

:

:- 2+ F

:

:

:

:-

MgF2

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Ionic Compounds - Properties

•crystalline•high melting point•high boiling point•soluble in water•electrolytes

•hard•brittle

Crystal Lattice

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Covalent Bonding - G.N. Lewis (1916)

Some atoms shareshare e- to form bonds. When two nonmetals bond, they often share electrons since they have similar attractions (EA) for them. This sharing of valence electrons is called the covalent bond.

Attraction Stable bond Repulsion

Number of bonds = Number shared e- pairs.

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Covalent Bonds

H .. :H H H+

H2

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Covalent Bonds

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Covalent Bonds - Epotential “well”

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single bond- one shared pair of e-

H − HHH

Lewis structures:

show ALL valence electrons

dot = 1 e- line = 1 pair of e-

Single Covalent Bonds

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# of e- sharedGroup # of to form an octet Example valence e- (8 - A group#)

4A 4 4 C in CH4

5A 5 3 N in NF3

6A 6 2 O in H2O

7A 7 1 F in HF H – F ....

Single Covalent Bonds..

H

H – C – H

H

H – O – H ....

H – F ....

F – N – F

F ..

..

......

..

...... ..

# of e- sharedto form an octet(8-A group#)

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Lewis Structures

Lewis electron-dot formulas or Lewis structures.

An electron pair is either a bonding pair (shared between two atoms) or a lone pair (an electron pair that is not shared).

bonding pair

lone pairs::H Cl

::

:H Cl

::bonding pair

lone pair:

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Multiple Bonds

In the molecules described so far, each of the bonds has been a single bond, that is, a covalent bond in which a single pair of electrons is shared.

It is possible to share more than one pair. A double bond involves the sharing of two pairs between atoms.

CC

H

H

H

H

orC:CH

H

H

H: : ::

:

C has octet.H OK with 2.

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Triple bonds are covalent bonds in which three pairs of electrons are shared between atoms.

Multiple Bonds

CC orHH

::

CC HH

:::

Elements that form multiple bonds: C, O , N, S

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The Procedure Using the molecular formula, count the total

number of valence electrons available (bonding + lone pairs). Valence electrons for each atom corresponds to group # Adjust for charge (add electron for each minus, delete

electron for each plus)

H2O

8 valence electrons

H3O+

8 valence electrons

2(1) + 6 3(1) + 6 - 1

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The Procedure Make a skeleton by connecting the atoms with single bonds

only. When connecting atoms, remember… Put the least electronegative atom in the center. (Usually the first listed

in the chemical formula.) Hydrogen is ALWAYS a terminal atom. More electronegative atoms

are terminal (F, O…) Make the structure symmetric.

H2O

8 valence electrons

H3O+

8 valence electrons

H O H H O H

H

4 electrons left over two electrons left over

4 pairs 4 pairs

2 pairs left 1 pair left

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The Procedure Put the left over electrons as lone pairs, preferably on

the more electronegative atoms Is the octet rule satisfied? If YES, then you’re done…

H2O

8 valence electrons

H3O+

8 valence electrons

H O H H O H

H

4 electrons left over two electrons left over

H O H H O H

H

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The Procedure

If you are electron-deficient (not enough electrons to complete an octet), then some atoms must share more than two electrons. “If you have a lone pair, make those two atoms share.”

Ex. C2H4

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The Procedure

If you have excess electrons, at least one atom must have an expanded valenceMust be element from third period or lowerUsually the central atom

e.g. SF4

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The total number of valence electrons on an atom (from bonds & lone pairs) cannot exceed that atom’s maximum valence. First period: 2 electrons (s) Second period: 8 electrons (s,p) Third period & below: prefer to have 8, but can expand

when necessary (s,p,d)

H O H H C H

H

H

P

Cl

Cl Cl

ClCl

General Rule

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Writing Lewis Dot Formulas

SCl220 e- total or 10 pairs

ClSCl

8 left

::

: :

::

::

0 left

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Cl

C

Cl

O

Writing Lewis Dot Formulas

COCl2

24 e- total 12 pairs

:

:: :

::

0 left

:

: :

Note that the carbon has only 6 electrons.

9 left

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Cl

C

Cl

O

Writing Lewis Dot Formulas

COCl2 12 pairs9 e- left

:

:: :

::

0 e- left

:: :

To fulfill the octet rule…“If you have a lone pair, make those two atoms share!”

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Writing Lewis Dot Formulas

COCl224 e- total

Cl

C

Cl

18 e- left

:

:: :

::

0 e- leftO: :

Note that the octet rule is now obeyed.

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Writing Lewis Dot FormulasPractice

N2 SF4

O2 ClO3-

HCN ClO2-1

PO4-3 NO3

-1

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Question

First evaluate the total valence electrons:24 e-

a. 26 e-, wrongb. 24 e-, looks OKc. 24 e-, one F has too manyd. 24 e-, N not enoughe. 24 e-, but least electronegative has to be in the centerf. 24 e-, no bond between two N.

no no

no no no

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Exceptions to the Octet Rule

Although many molecules obey the octet rule, there are exceptions where the central atom has less or more than eight electrons.

Incomplete octet – B, H

BF3Boron has 3 valence electrons

:F – B – F: .. .... ..

:F:..

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Exceptions to the Octet Rule

: F :

::: F :

F ::

:

: F

:: PF :

::

If a nonmetal is in the third period or greater it can accommodate as many as twelve electrons as the central atom.

PF5

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Exceptions to the Octet Rule

In sulfur tetrafluoride, SF4, the sulfur atom must accommodate two extra lone pairs for a total of 5 electron pair (10 electrons)

F ::

:

: F :

:

SF :

::

: F

:: :

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Formal Charge and Lewis StructuresIn certain instances, more than one feasible Lewis structure can be

illustrated for a molecule. For example,

H C N CNHor: :

The concept of “formal charge” can help us decide which structure is correct.

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Formal Charge and Lewis Structures

formal charge = valence e- before bonding– valence e- after bonding

= valence e- - [1/2 bonding e- + lone pair e-]

H C N CNHor: :

H: 1-½(2) = 0

C: 4 - ½(8) = 0

N: 5 – (½(8) + 2) = 0

H: 1-½(2) = 0

C: 4 – (½(6)+2) = -1

N: 5 – (½(8)) = +1

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Formal Charge and Lewis Structures

Smaller formal charges are more favorable More electronegative (or higher EA) atom should have

negative formal charges Like charges should not be on adjacent atoms Net formal charge should be the overall charge on the

molecule/ion.

orH C N:0 0 0

CNH :formal charges

0 +1 -1

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Practice

Determine the most stable structure for dinitrogen Determine the most stable structure for dinitrogen oxide. (All structures have 16 valence electrons.)oxide. (All structures have 16 valence electrons.)

N=N=ON=N=O N-N≡ON-N≡O N ≡ N - ON ≡ N - O-1 +1 0 -2 +1 +1 0 +1 -1

formal charge= valence e- - [1/2 bonding e- + lone pair e-]

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Practice - Formal Charge

Which structure is correct?

orO N Cl0 0 0

ClNO :-1 0 +1

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Delocalized Bonding: Resonance

The structure of ozone, O3, can be represented by two different Lewis electron-dot formulas.

O O

O:: :

::

:

OO

O

:::

::

:

The bond lengths for the above structures are:

O – O 132 pm O = O 112 pm

However, experiments show that both bonds are identical.

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Delocalized Bonding: Resonance

According to theory, one pair of bonding electrons is spread (delocalized) over the region of all three atoms.

In fact, the actual bond length is 127.8 pm (in between 132 and 112pm).

The actual molecule is a hybrid or composite structure and not different structures that change back and forth… although, we often represent it that way.

OO

O

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Delocalized Bonding: Resonance

Lewis resonance structures, have the same atoms in the same positions. Only an electron pair position is different.

OH O N

O

OH O N

O

OH O N

O

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Resonance Structures

O=S OO=S O O S=OO S=O

S=S OS=S O S O=SS O=S

Which pair does NOT represent resonance structures?

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Resonance StructuresDraw resonance structure(s) for the following:

O O C

O

-2

O O C

O

-2

O O C

O

-2

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“All covalent bonds are created equal but some are more equal than others.”

(We assumed equal sharing when we calculated formal charge.)

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Electroegativity vs Atomic Number

0

0.5

1

1.5

2

2.5

3

3.5

4

4.5

0 10 20 30 40 50 60

Atomic Number

Ele

ctro

neg

ativ

ity

(Pau

ling

)

He NeAr

KrXe

F

Cl Br

I

Electronegativity……is a measure of the ability of an atom in a molecule

to draw bonding electrons to itself when bonded.

decreasesdown a group

increasesacross a period

Periodic Trend - Electronegativity

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ElectronegativityNotice, there are NO values forEN for the noble gases.

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Types of Bonds

Ionic: ΔEN >1.8 electron transfer

Covalent: ΔEN <1.8 electron sharing

Metallic: electron-sea model or band theory

0.9 3.0

Na+ Cl-

2.5 2.1

C - H

1.6 1.6

Zn Zn

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Covalent Bonds (EN<1.8)

Non-polar covalent - ΔEN = 0 – 0.5

Examples: H-H, Cl-Cl, C-H bonds

Polar covalent - ΔEN = 0.5 – 1.8 Examples: H-O, C-Cl, C-O bonds

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46

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Bond Polarity

In HCl we have a partial negative charge on the chlorine (denoted -) and a partial positive charge on the hydrogen (denoted +)

The bond is polar covalent.

H :Cl::

:

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Bond PolarityArrange the following bonds from the most to the least polar:

HH, HCl, HF, HI, HBr

Compare the electronegativity of Cl, F, I and Br:Least Most Electronegative ElectronegativeI Br Cl F

Determine polarity:HH HI HBr HCl HFnon polar most polar

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PracticeWhich of the following bonds in each pair are more polar?

C-S or C-OCl-Cl or O=O N-H or C-H

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Bond LengthBond length (or bond distance) is the distance between the

nuclei of two bonded atoms (the sum of atomic radii).

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The size of atoms that form the bond determine the length.

C - N 147 pmC - C 154 pmC - P 187 pm (P, period 3)

Bond Length

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Bond Length – Multiple Bonds

As the electron density between atoms increases the bond lengths decrease; the atoms are pulled together more strongly.

decrease

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Bond EnthalpyBond Enthalpy – the enthalpy change that occurs when the bond

between two bonded atoms in the gas phase is broken and the atoms are separated completely at constant pressure.

As the electron density between two atoms increases, the bond

gets shorter and stronger.

Bond length Bond enthalpy

C

C

C

C

CC

154 pm

134 pm

120 pm 835 kJ/mol

602 kJ/mol

346 kJ/mol

NaCllattice energy-786 kJ/mol

MgOlattice energy-3791 kJ/mol

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Bond Enthalpy

Hº = ∑[(moles of bonds) × D(bonds broken)] -

∑[(moles of bonds) × D(bonds formed)]

Hº - standard enthalpy of reaction

Bond Enthalpies can be used to calculate the standard enthalpies of reaction (gas phase, STP)

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Bond Enthalpy

Estimate the Hº for the following reaction:

CH4(g) + 2O2(g) CO2(g) + 2H2O(g)

+ 2 :O = O: : O = C = O: + 2H – O - H

:: :: ::

4 C – H bonds 1 O = O bond 2 C = O bond 2 H – O bondper molecule per molecule per molecule per molecule

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Hº = ∑[(moles of bonds) × D(bonds broken)] -

∑[(moles of bonds) × D(bonds formed)]

+ 2 :O = O: : O = C = O: + 2H – O - H

:: :: ::

4 C – H bonds 1 O = O bond 2 C = O bonds 2 H – O bondsper molecule per molecule per molecule per molecule

= [4 × D(C-H) + 2 × D (O=O)] – [2 × D (C=O) + 4 × D(H-O)] =

[4 × 416 + 2 × 498] – [2 × 803 + 4 × 467] = -814 kJ

4 C – H bonds 2 O = O bonds 2 C = O bonds 4 H – O bonds