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Chemical Bonding I:Basic Concepts

Chapter 9

Valence electrons are the outer shell electrons of an atom. The valence electrons are the electrons thatparticipate in chemical bonding.

1A 1ns1

2A 2ns2

3A 3ns2np1

4A 4ns2np2

5A 5ns2np3

6A 6ns2np4

7A 7ns2np5

Group # of valence e-e- configuration Lewis Dot Symbol

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Lewis Symbols for Atoms

•Element symbol = nucleus + core electrons

•Valence electrons are drawn as dots around the symbol

•Up to 4 valence electrons are placed around the symbol one at a time; additional electrons are paired up

•The result is up to 4 pairs of electrons = octet

•NOTE: hydrogen can not have an octet. When forming bonds with other atoms, it can have a maximum of 2 electrons in its valence shell

O• •• • ••

Lewis Dot Symbols are not drawn for transition metals

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Li(s) + ½ F2(g) LiF(s)

1s22s1 1s2 = [He]Li Li+ + e-

1s22s22p5 1s22s22p6 = [Ne]

e- + F F - F -Li+ + Li+ F -

After F2 breaks apart into two neutral F atoms -

Lattice energy (E) increases as Q increases and/or as r decreases.

cmpd lattice energyMgF2

MgO

LiF

LiCl

2957

3938

1036

853

Q= +2,-1

Q= +2,-2

r F < r Cl

E = kQ+Q-r

Q+ is the charge on the cation

Q- is the charge on the anion

r is the distance between the ions

Lattice energy (E) is the energy required to completely separate one mole of a solid ionic compound into gaseous ions. It is always endothermic.

e.g. MgF2(s) → Mg2+(g) + 2F-(g)

Coulomb’s Law

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Born-Haber Cycle for Determining Lattice Energy

∆Hoverall = ∆Hrxn = ∆H1 + ∆H2 + ∆H3 + ∆H4 + ∆H5

solve for ∆H5 = -E

o ooooo

∆H5 < 0 exothermic

opposite to lattice energy E

∆H5 = -E

o

o

o

o

∆Hrxn = ∆H1 + ∆H2 + ∆H3 + ∆H4 + ∆H5o ooooo

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A covalent bond is a chemical bond in which two or more electrons are shared by two atoms, resulting in an octet for both atoms.

Lewis structure of F2

Lewis structure of H2O

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Double bond – two atoms share two pairs of electrons

Triple bond – two atoms share three pairs of electrons

N2

CO2

Guidelines for Drawing Lewis Structures(updated later on with the concept of “formal charge”)

1. Hydrogen is always a terminal atom because it canform only one bond.

2. The CENTRAL ATOM usually has the lowestelectron affinity (or electronegativity as defined later)

3. Arrange the atoms geometrically and symmetrically.

e.g. CH2O

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4. Sum up the total number of valence electrons (use thegroup number), and calculate the number of pairs.

5. Connect the atoms together so that each atom has anoctet (except H). You may have to form multiple bonds.

NF3

HNO3

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Lewis Structures of Charged Species

ClO-

NO2+

116C≡N

138C=N

143C-N120C≡C

133C=C

154C-C

Bond Length(pm)

Bond Type

Lengths of Covalent Bonds

Bond LengthsTriple bond < Double Bond < Single Bond

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Electronegativity is the ability of an atom to attract the electrons in a chemical bond towards itself.

Electron Affinity - measurable, Cl is highest

Electronegativity – Pauling Scale (relative scale), F is highest

X (g) + e- X-(g)

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H F FH

Polar covalent bond or polar bond is a covalent bond with greater electron density around one of the two atoms

electron richregionelectron poor

region e- riche- poor

δ+ δ-

Covalent

share e-

Polar Covalent

partial transfer of e-

Ionic

transfer e-

Increasing difference in electronegativity

Classification of bonds by difference in electronegativity

Difference Bond Type

0 Covalent≥ 2 Ionic

0 < and <2 Polar Covalent

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An atom’s formal charge is the difference between the number of valence electrons surrounding an isolated atom, and the number of electrons assigned to that atom in a Lewis structure.

N F

N O H

13

formal charge on an atom in a Lewis structure

=12

total number of bonding electrons( )

total number of valence electrons in the free atom

-total number of nonbonding electrons

-

The sum of the formal charges of the atoms in a molecule or ion must equal the charge on the molecule or ion.

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Formal Charge and Lewis Structures1. For neutral molecules, a Lewis structure in which there

are no formal charges is preferable to one in which formal charges are present.

2. Lewis structures with large formal charges are less plausible than those with small formal charges.

3. Among Lewis structures having similar distributions of formal charges, the most plausible structure is the one in which negative formal charges are placed on the more electronegative atoms.

What is the most likely Lewis structure for CH2O?

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A resonance structure is one of two or more Lewis structures for a single molecule that cannot be represented accurately by only one Lewis structure.

e.g. O3

e.g. what are the resonance structures of the carbonate (CO3

2-) ion?

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1. The Incomplete Octet (Central atom in Group 3A)

e.g. BF3

2. Odd-Electron Molecules

NO

NO2

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3. Expanded Octet – central atom has greater than 8 valence electrons surrounding it. Occurs only with elements in row 3 and higher because they have available d-orbitals

SF6

A. Covalent Molecules

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B. Polyatomic Ions

PO43- SO4

2-

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The enthalpy change required to break a particular bond in one mole of gaseous molecules is the bond energy.

H2 (g) H (g) + H (g) ∆H0 = 436.4 kJ

Cl2 (g) Cl (g)+ Cl (g) ∆H0 = 242.7 kJ

HCl (g) H (g) + Cl (g) ∆H0 = 431.9 kJ

O2 (g) O (g) + O (g) ∆H0 = 498.7 kJ O O

N2 (g) N (g) + N (g) ∆H0 = 941.4 kJ N N

Bond Energy

Bond Energies

Single bond < Double bond < Triple bond

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Average bond energy in polyatomic molecules

H2O (g) H (g)+ OH (g) ∆H0 = 502 kJ

OH (g) H (g)+ O (g) ∆H0 = 427 kJ

Average OH bond energy = 502 + 4272

= 464 kJ

Bond Energies (BE) and Enthalpy changes in reactions

∆H0 = total energy input – total energy released= ΣBE(reactants) – ΣBE(products)

Imagine reaction proceeding by breaking all bonds in the reactants and then using the gaseous atoms to form all the bonds in the products.

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Use bond energies to calculate the enthalpy change for:2C2H6(g) + 7O2(g) 4CO2(g) + 6H2O(g)

∆H0 = ΣBE(reactants) – ΣBE(products)

Type of bonds broken

Number of bonds broken

Bond energy (kJ/mol)

Energy change (kJ)

Type of bonds formed

Number of bonds formed

Bond energy (kJ/mol)

Energy change (kJ)