Basic Concepts of Chemical Bonding

Why are some compounds composed of ions and somecomposed of molecules?

The answer lies in the relationship between electronicstructure and chemical bonding, the forces holding atomstogether.

Ionic Bonds Covalent Bonds

Held together by electrostatic forces between particles of opposite charge

generally formed between metals and nonmetals

Formed by electron transfer

Formed by the sharing of electrons

Generally between two nonmetals of similar electronegativity

Metallic Bonds

Formed between atoms of the same metal

Freely moving electrons are delocalized around nuclei

Lustrous, good conductors of electricity

Ionic Bonds

Formed when a metal with low ionization energy combineswith a nonmetal of high electron affinity:

Na. + :Cl. Na+ + Cl-

Principal reason for ionic bond stability is the force of attraction between ions of unlike charge.

A measure of that stability is the lattice energy: energyrequired to completely separate a mole of a solid ioniccompound into its gaseous elements.

Magnitude of lattice energy defined by:Eel = k Q1Q2

d

Q1 and Q2 are the charges on the atoms measured in Coulombs; d represents the distance between the centersor size of the atoms, and k is a constant, 8.99 x 109 Jm/C2.

Lattice energy increases as the charge on the atoms increases and decreases as the atom gets larger.

Atoms should want to lose more electrons, Q1 and Q2 would be larger, but there is a balance between:

metals: lattice energy and ionization energy nonmetals: lattice energy and electron affinity

Transition metals lose valence shell s electrons first and then as many d electrons as need to form an ion of a particular charge. “s before d”

Arrange the following compounds in order of increasinglattice energy:

NaF CsI CaO

The lattice energies for KF, MgO, and ScN are as follows:808, 3795, and 7547 kJ/mol respectively. Account for the trend in lattice energy.

Born-Haber Cycle

Hess’s Law like series of steps that allows us to calculateHf of a compound.

Hfo = Na(s) + 1/2 Cl2(g) NaCl(s) -410.9 kJ/mol

Step 1: Generate gaseous atoms of Na and ClNa(s) Na(g) Hf

o = 107.7 1/2 Cl2(g) Cl(g) 121.7

Both processes are endothermic

Step 2: Remove an electron from Na and add it to ClNa(g) Na+(g) + 1e- I1(Na) = 496 kJ/molCl(g) + 1e- Cl-(g) Ea = -349

Step 3: Bring the gaseous ions together to form NaCl(s) Na+(g) + Cl-(g) NaCl(s) H = ?

Add the steps together.

Determine the Born-Haber cycle for the formation of CaBr2.

Ionic Size

Crucial to structure and stability of ionic solids

size determines packing and lattice energyCations are smaller than parent atoms, why?Anions are larger than parent atoms, why?

Arrange in order of decreasing size: Mg2+, Ca2+, Ca.

Arrange the following in order of decreasing size:

S2-

K+

Ca2+

Cl-

Sr2+

Y3+

Rb+

Covalent Bonding

Forms compounds that are gases, liquids, and soft solids.

Low melting points, vaporize easilyShare a pair of electronsElectron density concentrated between nuclei.Lewis electron dot structures are used to

represent the bonding arrangement in an ion or molecule.

Lewis StructuresUse NAS formula for molecules obeying octet rule.

N = number of electrons needed; multiply all non-H atoms by 8; H by 2.A = number of electrons available; add up valence electrons.S = N-A represents the shared electrons. Divide this number by 2 = number of bonds that must be drawn.

Draw the Lewis structure for the following:NF3 CH3Cl ClO4

- CO2

Bond Polarity and Electronegativity

When a covalent bond is formed, the electron pair may not be equally shared by both atoms; in this case, the bond is said to be polar.

A non-polar covalent bond is one where the atoms are shared equally.

Electronegativity is used to estimate the polarity of a bond; EN is the ability of an atom to attract electrons to itself.

Pauling’s electronegativity scale was developed to quantifyEN based on thermochemical data. On Pauling’s scale, F isthe most EN element and has an assigned value of 4.0, Cshas the lowest value, 0.7. The trend for EN is:

I

The difference in EN can be used to determine bond polarity. EN difference 0.0-0.4 non-polar covalent;0.4-1.7 polar covalent; >1.7 ionic

Which bond is more polar?

B - Cl C - Cl

P - F P - Cl

Dipole Moments

A polar molecule is one in which the centers of positive and negative charge do not coincide.

Molecules with polar bonds that are not symmetric are said to be polar and have a dipole moment > 0.

A molecule’s polarity helps determine many of its physical characteristics including state.

Dipole Moments

While it is not necessary to calculate dipole moments, you should be aware of the following:

symbol used: = Qrmeasured in Debye (D)higher the value, more polar the molecule

Molecular nomenclature:Molecules are named by exactly describing the

number of atoms present using a prefix system.

OF2 = oxygen difluoride

Formal Charge

-Means to decide between what appear to be equivalentLewis structures.

Formal charge = Number of valence electrons - (1/2 thenumber of bonding electrons + number of non-bonding electrons)

Choose structure with smallest formal charge.

NO2- N = 5- ((1/2)(6) + 2) = 0

O = 6- ((1/2)(4) + 4) = 0 O = 6- ((1/2)(2) + 6) = -1

Draw Lewis structures for the following molecules,determine formal charge and select the most likelystructure in each case:

NNO or NON

HCN or HNC

NOBr or ONBr

http://cheminf.cmbi.ru.nl/wetche/organic/

Resonance Structures

Two alternative Lewis structures that are equivalent.

All forms are equivalent but only one form of the molecule exists that is observed experimentally.

Lewis structures are limited in describing electron distribution.

Exceptions to the Octet Rule

NAS cannot be used here.Molecules with an odd number of electrons: ClO2,

NOMolecules in which an atom has less than an octet:

BF3

Molecules in which an atom has more than an octet: PCl5, SF6. Second period elements will NEVER expand their octets as no d electrons are available.

Strengths of Covalent Bonds

Stability of molecule relates to the strengths of the covalent bonds.

Use average bond energies to calculate enthalpy of reaction.

H = Bonds Broken - Bonds FormedMolecules with strong chemical bonds have

less tendency to undergo chemical change.

Using the table of Bond enthalpies, calculate the enthalpyof reaction for the combustion of ethane.

Bonds Lengths and Bond Strengthsgenerally as the number of bonds increases, bondlength decreases and bond gets stronger.

Application of bond stability: explosivesdecompose exothermicallyproduce large amounts of gasdecompose rapidlyusually produce N2, CO2 or CO