Calcium carbonate is one of the most in-teresting and versa-tile compounds on the planet.

Roughly 4% of earth’s crust

The major compo-nent of rocks such as limestone and mar-ble

The white cliffs of Dover are one of the most famous natural formations of CaCO3.

THE PROPERTIES OF SUBSTANCES &

CHEMICAL BONDS The properties of substances are determined in large

part by the chemical bonds (O2 vs N2)

What determines the type of bonding in each substance? The electronic structures of the atoms are the key to an-

swering the question

We will examine the relationship between the electronic structures of atoms and the chemical bonds they form

8.1 CHEMICAL BONDS, LEWIS SYMBOLS, AND THE

OCTET RULE Three general types of chemical

bonds•Ionic bond: electrostatic forces•Covalent bond: the sharing of

electrons•Metallic bond: free electrons

The valence electrons•The electrons involved in chemical bonding•In most atoms, they reside in the outermost occupied shell

The Lewis symbol•Diagrams that show the bonding between atoms of a mole-

cule and the lone pairs of electrons that may exist in the mole-cule

8.1 CHEMICAL BONDS, LEWIS SYMBOLS, AND THE OCTET RULE

LEWIS SYMBOLS

The noble gases have very stable electron arrangements

(high ionization energies)

Atoms often gain, lose, or share electrons to achieve the

same number of electrons as the noble gas

The octet rule

•Atoms tend to gain, lose or share electrons until they are sur-

rounded by eight valence electrons

•The rule provides a useful framework to introduce concepts

of bonding although there are many exceptions

8.1 CHEMICAL BONDS, LEWIS SYMBOLS, AND THE OCTET RULE

THE OCTET RULE

8.2 IONIC BONDING

The reaction can be explained by the low ionization energy of Na and the high electron affinity of Cl

We can also explain the reaction by the octet rule

Consider a reaction of Na(s) and Cl2(g)

8.2 IONIC BONDING

FORMATION OF SODIUM CHLORIDE

The formation of sodium chloride is very exothermic

The heat of formation of other ionic substances is also quite

negative

The first ionization energy of Na(g) is 496 kJ/mol

The electron affinity of Cl(g) is −349 kJ/mol

If we consider only the electron transfer for the formation reac-

tion, the energy change would be:

496 − 349 = 147 kJ/mol

How can we explain this?

8.2 IONIC BONDINGENERGETICS OF IONIC BOND FOR-

MATION

The lattice energy is the energy required to completely sepa-rate a mole of a solid ionic compound into its gaseous ions

The potential energy of two interacting charged particles

The magnitude of lattice energies depends predominantly on the ionic charges

8.2 IONIC BONDING

THE LATTICE ENERGY

8.2 IONIC BONDING

THE LATTICE ENERGY

Sample Exercise 8.1 Magnitudes of Lattice Energies

Without consulting Table 8.2, arrange the following ionic compounds in order

of increasing lattice energy: NaF, CsI, and CaO.

Answer: CsI < NaF < CaO.

Which substance would you expect to have the greatest lattice energy, MgF2,

CaF2, or ZrO2?

Practice Exercise

Answer: ZrO2

8.2 IONIC BONDING

The lattice energy can not be determined directly by experi-ment

The formation of Na2+ and Cl2− are energetically very unfavorable Groups 1A, 2A, and 3A atoms form 1+, 2+, and 3+ ions, respectively Groups 5A, 6A, and 7A atoms form 1−, 2−, and 3− ions, respectively Transition metals do not observe the octet rule: Fe2+ and Fe3+

8.2 IONIC BONDINGELECTRON CONFIGURATIONS OF

IONS

A chemical bond formed by sharing a pair of electrons

The attractions and repulsions among electrons and nuclei in the H2 molecule

Quantum mechanical calcula-tion tells us that the concentra-tion of electron density between the nuclei leads to a net attrac-tive force that constitutes the covalent bond holding the mol-ecule together

8.3 COVALENT BONDING

The formation of covalent bonds can be represented using Lewis symbols

8.3 COVALENT BONDING

LEWIS STRUCTURES

For the nonmetals, the number of valence electrons in a neu-tral atom is the same as the group number

the number ofcovalent bonds 1 2 3 4

The Lewis structures of CO2 and N2

8.3 COVALENT BONDING

MULTIPLE BONDS

Bond length: the distance between the nuclei of atoms in-volved in a bond

Compare the electron-sharing of Cl2 and H2 with the shar-

ing of H2O and NaCl

Polar covalent bond and nonpolar covalent bond If the difference of the covalent bond-forming atoms in rela-

tive ability to attract electrons is large enough, an ionic bond is formed

Electronegativity is defined as the ability of an atom in a molecule to attract electrons to itself

8.4 BOND POLARITY AND ELECTRONEGA-TIVITY

Linus Pauling (1901-1994) developed the first and most widely used electronegativity scale

Electronegativity can be used to estimate whether a given bond will be nonpolar or polar covalent, or ionic

The electronegativity of an atom in a molecule is related to its ionization energy and electron affinity

Fluorine, the most electronegative element, has an elec-tronegativity of 4.0

The least electronegative element, cesium, has an elec-tronegativity of 0.7

8.4 BOND POLARITY AND ELECTRONEGATIVITY

ELECTRONEGATIVITY

8.4 BOND POLARITY AND ELECTRONEGATIVITY

ELECTRONEGATIVITY

The greater the difference in electronegativity between two atoms, the more polar their bond

8.4 BOND POLARITY AND ELECTRONEGATIVITYELECTRONEGATIVITY AND BOND

POLARITY

8.4 BOND POLARITY AND ELECTRONEGATIVITYELECTRONEGATIVITY AND BOND

POLARITY

We can indicate the polarity of the HF molecule in two ways:

8.4 BOND POLARITY AND ELECTRONEGATIVITY

DIPOLE MOMENTS

Polarity helps determine many of the properties of substances such as hydrogen bonding and solvation

How can we quantify the polarity of molecule? For the two equal and opposite charges, the dipole moment is:

• , dipole moment (C·m or debye, D)

• Q, charge (C)• r, distance (m)

1 D = 3.34 × 10-30 C·m

Sample Exercise 8.5 Dipole Moments of Diatomic Molecules

The bond length in the HCl molecule is 1.27 Å. (a) Calculate the dipole mo-ment, in debyes, that would result if the charges on the H and Cl atoms were 1+ and 1–, respectively. (b) The experimentally measured dipole moment of HCl(g) is 1.08 D. What magnitude of charge, in units of e, on the H and Cl atoms would lead to this dipole moment?

Solution

The actual charges on the atoms decrease from 0.41 in HF to 0.057 in HI

8.4 BOND POLARITY AND ELECTRONEGATIVITY

DIPOLE MOMENTS

The ability to quickly categorize the predominant bonding in-teractions in a substance as covalent or ionic imparts consid-erable insight into the properties of that substance

By considering the interaction between a metal and a non-metal

•SnCl4: colorless liquid, mp -33 ˚C, bp 114 ˚C, polar covalent

To use the difference in electronegativity

•SnCl4: 1.2; NaCl: 2.1

•MnO: 2.0, ionic; Mn2O7: 2.0, polar covalent

•The increase in the oxidation state of a metal leads to an in-crease in the degree of covalent character in the bonding

8.4 BOND POLARITY AND ELECTRONEGATIVITYDIFFERENTIATING IONIC AND COVALENT

BONDING

1. Find the sum of valence electrons of all atoms in the polyatomic ion or molecule.– If it is an anion, add one

electron for each negative charge.

– If it is a cation, subtract one electron for each posi-tive charge.

PCl35 + 3(7) = 26

8.5 DRAWING LEWIS STRUCTURES Lewis structures can help us understand the bonding in many

compounds and are frequently used when discussing the properties of molecules

2. Arrange atoms. The central atom is the least electronegative el-ement that isn’t hydro-gen. Connect the outer atoms to it by single bonds.

Keep track of the electrons: 26 - 6 = 20

8.5 DRAWING LEWIS STRUCTURES

3. Complete the octets around all the atoms bonded to the central atom. Hydrogen is an exception

Keep track of the electrons: 26 - 6 = 20; 20 - 18 = 2

8.5 DRAWING LEWIS STRUCTURES

4. Place any leftover electrons on the cen-tral atom

Keep track of the electrons:26 - 6 = 20; 20 - 18 = 2; 2 - 2 = 0

8.5 DRAWING LEWIS STRUCTURES

5. If there are not enough elec-trons to give the central atom an octet, try multiple bonds

8.5 DRAWING LEWIS STRUCTURES

Sample Exercise 8.8 Lewis Structure for a Polyatomic Ion

Practice ExerciseDraw the Lewis structure for (a) ClO2

–, (b) PO43–

Draw the Lewis structure for the BrO3– ion.

SolutionThe total number of valence electrons is, therefore, 7 + (3 6) + 1 = 26. For oxyanions— BrO3

–, SO42–, NO3

–, CO32–, and so forth—the oxygen atoms surround the central non-

metal atom.

Answers: (a) (b)

The charge the atom would have if all the atoms in the molecule had the same electronegativity

How to assign formal charges.•For each atom, count the electrons in lone pairs and

half the electrons it shares with other atoms.•Subtract that from the number of valence electrons for

that atom: the difference is its formal charge.

8.5 DRAWING LEWIS STRUCTURE

FORMAL CHARGE

The concept of formal charge helps us choose a pre-ferred Lewis structure

How to choose the correct structure.•Choose the Lewis structure in which the atoms bear

formal charges closest to zero•Choose the Lewis structure in which any negative

charges reside on the more electronegative atoms

FORMAL CHARGE8.5 DRAWING LEWIS STRUCTURE

preferred

FORMAL CHARGE8.5 DRAWING LEWIS STRUCTURE

preferred

Oxidation number Formal charge Actual partial charge

+1 −1 0 0 +0.18 −0.18

+0.178 −0.178

Dipole moment

8.5 DRAWING LEWIS STRUCTURE

Consider the Lewis structure of ozone, O3

8.6 RESONANCE STRUCTURES

Both Lewis structures are not corre-sponding to the structure of ozone

Resonance structure: An alternate way of drawing a Lewis dot structure for a compound

How can we understand the resonance structure? The rules for drawing Lewis structures do not allow us to

have a single structure that adequately represents the ozone molecule

8.6 RESONANCE STRUCTURES

8.6 RESONANCE STRUCTURES

Draw the resonance structures for HCO2− and NO3

− .

8.6 RESONANCE STRUCTURES

RESONANCE IN BENZENE

All six C-C bonds are of equal length 1.40 ÅC-C single bond 1.54 Å; C=C double bond 1.34 Å

8.7 EXCEPTIONS TO THE OCTET RULE

ODD NUMBER OF ELECTIONS In a few molecules and polyatomic ions, such as ClO2, NO,

NO2, and O2−, the number of valence electrons is odd

Complete pairing of these electrons is impossible

An octet around each atom cannot be achieved

NO contains 5 + 6 =11 valence electrons

8.7 EXCEPTIONS TO THE OCTET RULELESS THAN AN OCTET OF VALENCE ELECTRONS Consider boron trifluoride, BF3

In the Lewis structure of BF3, there are only six electrons around the boron atom and the for-mal charges are zero

If you try to complete the octet around boron: Most important

8.7 EXCEPTIONS TO THE OCTET RULEMORE THAN AN OCTET OF VALENCE ELECTRONS

Consider phosphorus pentachloride, PCl5

Elements from the third period and beyond have unfilled nd orbitals that can be used in bonding

The larger the central atom is, the larger the number of atoms that can surround it

Draw the Lewis structure of SF4, ICl4−, and PO43−

The orbital diagram for the valence shell of a phosphorus atom

8.8 STRENGTH OF COVALENT BONDS

The stability of a molecule is related to the strengths of the covalent bonds it contains

The strength of a bond is measured by determining how much energy is required to break the bond

The bond enthalpy for a Cl-Cl bond, D(Cl-Cl), is measured to be 242 kJ/mol

8.8 STRENGTH OF COVALENT BONDS

8.8 STRENGTH OF COVALENT BONDS

The bond enthalpy is always positive as energy is always re-quired to break chemical bonds

A molecule with strong chemical bonds generally has less tendency to undergo chemical change than does with weak bonds

8.8 STRENGTH OF COVALENT BONDSBOND ENTHALPIES AND THE EN-

THALPIES OF REACTION We can use average bond enthalpies to estimate the en-

thalpies of reactions in which bonds are broken and new bonds are formed

Consider a gas-phase reaction:

CH4(g) + Cl2(g) → CH3Cl(g) + HCl(g)

8.8 STRENGTH OF COVALENT BONDSBOND ENTHALPIES AND THE EN-

THALPIES OF REACTION

8.8 STRENGTH OF COVALENT BONDSBOND ENTHALPY AND BOND LENGTH

Enormous amounts of energy can be stored in chemical bonds → can be used as an explosives

Characteristics of an explosive•Very exothermic decomposition•Gaseous products•Rapid decomposition•Controllably stable

− 2− − ++ ++