Chapter 9Chemical Bonding I: The

Lewis Model

Gang Chen, Department of Chemistry, UCF

Lecture 1&2 Review

2

Chemical Bonding

• All bonds are created by the exchange or sharing of

electrons.

• The exchange or sharing of electrons results in lower

energy for the compound relative to the separate atoms.

• Ionic: exchange; cations and anions present

• Covalent: sharing between atoms

• Metallic: sharing by forming a mobile “sea of electrons”

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The Ionic Bond

19K 1s22s22p63s23p64s1 ⎯→ 19K+ 1s22s22p63s23p6 + e-

• When a metal atom loses electrons it becomes a cation.

• Metals have low ionization energy, making it relatively

easy to remove electrons from them.

• When a nonmetal atom gains electrons it becomes an anion.

• Nonmetals have high electron affinities, making it

advantageous to add electrons to these atoms.

• The oppositely charged ions are then attracted to each other,

resulting in an ionic bond.

9F [He] 2s22p5 + e- ⎯→ 9F- [He] 2s22p6

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The Covalent Bond

• A covalent bond is based on the sharing of pairs of electrons

between two atoms.

• Nonmetal atoms have relatively high ionization energies, so it is difficult to remove electrons from them.

• Covalent bonding lowers energy by sharing electrons between two nonmetals.

• Shared electrons hold the atoms together by attracting nuclei of both atoms.

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Octet Rule

• Valence electrons (outermost) are involved in forming bonds.

• Octet rule - an atom will form chemical bonds to achieve a

complement of eight valence electrons.

• During ionic bond formation, the cations and anions achieve

ns2np6 electronic configurations (noble gas configuration).

• Metals lose electrons.

• Nonmetals gain electrons.

• During covalent bond formation, electrons are shared

between two atoms.

• Shared electrons are available to both bonding atoms.

• Sharing leads to 8 valence electrons around each atom.

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Lewis Structures of Atoms

• Lewis dot symbols keep track of valence electrons,

especially for main group elements, allowing prediction of

bonding in molecules.

• To draw a Lewis dot symbol, the valence electrons are

represented by dots and are placed around the element

symbol.

• The first four dots are placed singly.

• Starting with the fifth dot, they are paired.

• The second period Lewis symbols are shown below:

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Lewis Structures of Atoms

• Lewis dot symbols for main group elements.

• Elements within a group have the same number of

valence electrons and identical Lewis dot symbols.

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Lewis Structures of Ions

• Lewis dot symbols can be used to represent the ions formed

by atoms.

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Lattice Energy

• Forming the ion pair from a metal and nonmetal usually requires energy input.

• Ionization energies are positive.

• Electron affinities for nonmetals are negative.

• Energy input to form the cation is not offset by energy released by forming the anion.

Na(g) → Na+(g) +496 kJ

Cl (g) → Cl−(g) −349 kJ

Na(g) + Cl(g) → Na+(g) + Cl−(g) +147 kJ

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Lattice Energy

• The crystal lattice formed due to the electrostatic force of attraction between the ion pairs significantly lowers overall energy (energy is released when the ionic bond forms).

• Electrostatic attraction is nondirectional, therefore, there is no ionic molecule

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Lattice Energy

• The lattice energy is the energy released when the solid crystal forms from separate ions in the gas state.

• The potential energy, V, for the ion pair can be calculated.

• k = 1.389 x 105 kJ pm/mol

V = kq1q2

r

F µq1q2

r2

• Coulomb’s Law: F is force between two ions, q1 and q2 are the charges, and r is the distance between the nuclei of the two ions.

Na(g) + Cl(g) → Na+(g) + Cl−(g) +147 kJ

Na+ (g) + Cl−(g) → NaCl(s) −788 kJ

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Lattice Energy

• From the potential energy equation, we can see that:

• Small ions with large charges form ionic compounds with

large lattice energies.

• Large ions with small charges form ionic compounds with

small lattice energies.

V = kq1q2

r

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Example ProblemIn each of the following pairs of ionic solids, the ions are

arranged into similar lattices. Which one has the larger

lattice energy?

b) NaF or CsI

a) CaO or KF CaO larger charges; 22

r

qqkV 21 −+

=

NaF smaller ions; r

Small ions with large charges form ionic compounds with

large lattice energies.

Large ions with small charges form ionic compounds with

small lattice energies.

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Ionic Bonding: Model versus Reality

• Ionic compounds have high melting points and boiling

points.

• The attractions between ions are strong.

• The stronger the attraction (larger the lattice energy),

the higher the melting point.

• The solid state of ionic compound does not conduct

electricity.

• The liquid state and the solution of ionic compounds

conduct electricity.

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The Covalent Bond

• Lewis dot structures show how electrons are shared in a

molecule.

• A pair of shared electrons between two atoms is a bonding

pair.

• Bonding pairs represented by a line between two atomic

symbols.

• Pairs of electrons associated with one atom are

nonbonding or lone pair electrons.

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The Covalent Bond

• By sharing an electron from each atom, two hydrogen atoms

can form a covalent bond.

• Hydrogen violates the octet rule by sharing only two

electrons (duet).

• When two fluorine atoms combine, they form a stable

covalent bond.

• By sharing a pair of electrons, each atom is surrounded

by eight valence electrons.

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The Covalent Bond

• When two atoms share one pair of electrons, it is called a single covalent bond.

• When two atoms share two pairs of electrons the result is called a double covalent bond.

• When two atoms share three pairs of electrons the result is called a triple covalent bond.

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Covalent Bonding: Model versus Reality

• Lewis theory implies that some combinations should be stable, whereas others should not.• Because the stable combinations result in “octets”

• Compounds of nonmetals are made of individual molecule units.• the attractions between covalently bonded atoms are

directional.

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Covalent Bonding: Model versus Reality

• Molecular compounds have low melting points and boiling points.• The covalent bonds are strong, but the attractions between

the molecules are generally weak.

• Molecular compounds do not conduct electricity in the solid or liquid state.• There are no charged particles around to allow the material

to conduct.