chapter 8: bonding: general concepts chapter 8: bonding: general concepts cartoon courtesy of...

Chapter 8:

Bonding:Genera

l Concept

s

Cartoon courtesy of NearingZero.net

Chemical Bonding Forces that hold groups of atoms together and make them function as a unit.

Bond EnergyEnergy required to break a bond

Bond Polarity and Dipole Moments

Dipolar Molecules1. Molecules with a somewhat negative end and a somewhat positive end (a dipole moment)2. Molecules with preferential orientation in an electric field

+ + + + + + + +

- - - - - - - -3. All diatomic molecules with a polar covalent bond are dipolar

Bond Polarity and Dipole Moments

Molecules with Polar Bonds but no Dipole Moment1. Linear, radial or tetrahedral symmetry of charge distribution

a. CO2 - linearb. CCl4 – tetrahedral

See table 8.2

Ionic Bonding

Ionic bond: the electrostatic force that holds ions together in an ionic compound.

Examples of Ionic Compounds (aka Salts):NaClBaCl2

Ionic Bonding Electrons are transferred

Electronegativity differences are

generally greater than 1.7 The formation of ionic bonds is

always exothermic!

Determination of Ionic

Character

Compounds are ionic if they conduct electricity in their molten state

Electronegativity difference is not the final determination of ionic character

Properties of Ionic Compounds

Structure: Crystalline solids

Melting point:

Generally high

Boiling Point:

Generally high

Electrical Conductivity:

Excellent conductors, molten and aqueous

Solubility in water:

Generally soluble

Coulomb’s Law

r

QQnmJxE 2119 )1031.2(

“The energy of interaction between a pair of ions

• E is in Joules• r is the distance between the center of the ions

• Q1 and Q2 are the charges of the ions• A negative quantity indicates attraction • A positive quantity indicates repulsion

Coulomb’s Law

• Example: In solid NaCl the distance between the centers of the ions is 2.76 Å (0.276 nm) Calculate the ionic energy per pair of ions:

Sodium Chloride Crystal Lattice

Ionic compounds form solid crystals at ordinary temperatures.

Ionic compounds organize in a characteristic crystal lattice of alternating positive and negative ions.

All salts are ionic compounds and form crystals.

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

Electronegativity - F is highest

X (g) + e- X-

(g)

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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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Classify the following bonds as ionic, polar covalent, or covalent: The bond in CsCl; the bond in H2S; and the NN bond in H2NNH2.

Cs – 0.7 Cl – 3.0 3.0 – 0.7 = 2.3Ionic

H – 2.1 S – 2.5 2.5 – 2.1 = 0.4Polar Covalent

N – 3.0 N – 3.0 3.0 – 3.0 = 0 Covalent

Formation of Ionic compounds

• Stable compounds are formed when nonmetallic elements take electrons from metals.

• Atoms usually have a noble gas configuration

Formation of Ionic compounds

In general: • When a binary ionic compound is

formed – the nonmetal has noble gas

configuration– The valence orbitals of the

representative metal is emptied

• The term ionic compounds refers to the solid state of the compound

• A collection of positive and negative ions arranged to minimize repulsions and maximize attractions

Predicting formulas of ionic compounds

Predicting formulas of ionic compounds

• Large electronegativity differences between atoms mean electrons will be transferred

Predicting formulas of ionic compounds

• Hydrogen typically behaves as a nonmetal

• The number of electrons transferred depends on how many each atom needs to gain or lose to achieve noble gas notation

Predicting formulas of ionic compounds• EXCEPTIONS:

– Tin forms Sn2+ and Sn4+

– Lead forms Pb2+ and Pb4+

– Bismuth forms Bi3+ and Bi5+

– Thallium forms Tl+ and Tl3+

Energy and Binary Ionic Compounds

• Factors that influence stability and structure

• Ionic compounds form because together they have lower energy than the original elements

Lattice Energy

• The energy released when an ionic solid is formed from its ions

• LE is negative (exothermic)• Used as a step to calculate

energy of formation

MX(s)(g)X(g)M

Lattice energy increases as Q increases and/or as r decreases.

CompoundLattice Energy (kJ/mol)

MgF2

MgO

LiFLiCl

29573938

1036853

Q: +2,-1Q: +2,-2

r F- < r Cl-

Q+ and Q- is the charge on the cation and anionr is the distance between the ions

E is the potential energy

k is a constant based on the compound

r

QQkE

Calculating Energy of formation Hf

• If we know the steps in the process then we can apply Hess’s law

• Because energy is a state function

• Break the reaction up into steps• Add them up

Estimate Hf for Sodium Chloride

Na(s) + ½ Cl2(g) NaCl(s)Lattice Energy -786 kJ/mol

Ionization Energy for Na 495 kJ/mol

Electron Affinity for Cl -349 kJ/mol

Bond energy of Cl2 239 kJ/mol

Enthalpy of sublimation for Na

109 kJ/mol

Na(s) Na(g) + 109 kJNa(g) Na+(g) + e- + 495 kJ

½ Cl2(g) Cl(g) + ½(239 kJ)Cl(g) + e- Cl-(g) - 349 kJ

Na+(g) + Cl-(g) NaCl(s) -786 kJ Na(s) + ½ Cl2(g) NaCl(s) -412

kJ/mol

The energy diagram for the formation of MgO and NaCl.

The Lattice energy to combine Mg2+ and O2- is much more negative than the energy needed for the process that produces Mg2+ and O2- ions.