Chapter 8 Bonding

8.1 Types of Bonds• There are lots of experiments we can do

to determine the nature of materials– Melting point– Conductivity– Solubility– Charge distribution in an electric field

• Bond Energy – The energy required to break a bond– Tells us the strength of a bonds interactions.

Ionic Compounds• We all know that ionic bonding is a

result of electrostatic attractions of oppositely charged ions.

• And ionic compounds are formed when a nonmetal and a metal react.

Coulomb’s law• Coulomb’s law is used to measure the

energy of interaction between a pair of ions.

• Where E is energy in joules, r is the distance between the ion centers in nm, and Q1 and Q2 are the charges of the ions.

19 1 2(2.31 10 )QQ

E x J nmr

Coulomb’s law in use• Let’s look at salt with ions

at the given distance:

• Notice the E is negative (indicating an attractive force) which means the ion pair has LOWER energy than the separated ions.

19 1 1(2.31 10 )

0.276E x J nm

nm

198.37 10E x J

Repulsive forces• Let’s look at H-H

bonds. When will a H2 molecule be favored

• A bond will form if the energy of the aggregate (whole created from its parts) is lower than that of the separated ions!!

An energy profile

Bond length• The distance where the energy is

MINIMAL is called the bond length.

• The type of bond we seen in an H2 molecule is called a covalent bond. (where electrons are shared) A mutual attraction of the two nuclei for the shared electrons.

\

Polar Covalent Bonds• In between the two extreme types of

bonding (ionic and covalent) is an intermediate case; called polar covalent bonds.

• These bonds are between atoms that are not so different that electrons are completely transferred but are different enough that there is unequal sharing.

H—F • Look at the HF molecule below. The

symbol δ (lowercase delta) indicates a fractional charge.

• This indicates that fluorine has a stronger attraction for the shared e- than hydrogen does.

H F

electron richregion

electron poorregion

8.2 Electronegativity• Electronegativity – is the ability of an

atom in a molecule to attract shared electrons to itself.

• Linus Pauling’s model is what we used to assign a value to electronegativity.

• The trend is that electronegativity generally increases across a period and decreases down a group.

Relationship between electronegativity and bond

type.

Using electronegativity to determine bond

polarity• Using electronegativity values, arrange the following bonds in order of increasing polarity: H – H, O – H, Cl – H, S – H, and F – H.

Remember bond polarity is the difference of the electronegativities of the atoms

forming the bond

The answer!!In increasing polarity

Bond Difference Bond typeH – H 0 Covalent

S – H 0.4 Polar covalent

Cl – H 0.9 Polar covalentO – H 1.4 Polar covalentF – H 1.9 Polar

covalent

P

O

L

A

R

I

T

Y

8.3 Bond Polarity & Dipole moments

• When a molecule has a center of positive charge and center of negative charge it is said to be dipolar or have a dipole moment.

• The dipolar character of a molecule is often represented by an arrow that points toward the negative charge center and the tail indicates the positive center of charge.

Another way to look at it…

• Electrostatic potential maps also show the charge distribution of a molar molecule.

• Red shows the electron rich area and the violet shows theelectron poor region.

• Any diatomic molecule with a polar bond will show a molecular dipole moment

Sometimes…• Polyatomic molecules exhibit dipolar

behavior. And in an electric field, water acts as if it has 2 centers of charge.

Other times…• Individual bond polarities are

arranged in such a way that they “cancel” each other out. As seen below in the CO2 molecule.

• Polar bonds but NO dipole moment.

• There are many more of these cases!!

Types of Molecules with Polar Bonds but No

Resulting Dipole Moment

Try these.• Show the direction of the bond polarities

and indicate which ones have a dipole moment:

• HCl• Cl2

• SO3

• CH4

• H2S

Cl2 has NO bond polarity because the e- are shared equally. Therefore, there is NO dipole moment.

8.4 Ions: Electron Configuration & Sizes

• Generalizations:– When 2 nonmetals react to form a

covalent bond, they share e- so that the valence electron configuration of both atoms is complete. (i.e. both nonmetals have attained noble gas electron configurations)

– When a nonmetal and a representative metal react to form a binary ionic compound, ions form so that the valence e- configuration of the nonmetal achieves the e- config of the next noble gas and the valence orbitals of the metal are emptied. (both ions attain noble gas e- configurations)

Predicting formulas of ionic compounds

• When the term ionic compound is used, typically, the state of matter being referred to is a SOLID.

• The + and – ions are packed together in such a way so thatthe + + and - - repulsions are minimized and the + - attractions are maximized.

Sizes of Ions• The size of an ion plays an

important role in the stability of an ionic solid and the properties of the ions in solution.

• Most of the time ionic radii are determined by measuring the distances between ions in ionic compounds… BUT this assumes how the distance is divided between the two ions.

Controversy• This method of determining ion

size creates a controversy…sooo….we will concentrate on the trends!

• Consider relative ion size and the size of the parent atom. + ions are formed by removing an e- so the resulting cation is SMALLER than its parent atom. And the opposite is true for anions.

Depends on parents• Remember the generalizations we

talked about…

• Size depends on the “parent’s” position in the periodic table. A given period has both elements that give up and gain electrons. So some elements get larger (to emulate the next noble gas) and some elements get smaller (to “look like” the previous noble gas)

Isoelectronic ions.• Isoelectronic ions are ions that

contain the same number of electrons.– For example O2-, F-, Na+, Mg2+, and

Al3+ each have 10 electrons ([Ne]) so the amount of repulsions should be the same.

• Let’s look at how their sizes vary.

What other factor should we look at for size?

• When looking at isoelectronic ions, the other factor to consider is the number of protons in the nucleus. The number of protons increases from 8 to 13 from O2- to Al3+.

• So as the number of protons increases, the ATTRACTION to the 10 electrons is GREATER and it causes the ions to become smaller.

8.5 Energy Effect in Binary Ionic Compounds

• We know that metals and nonmetals react by transferring e- and that the result is a solid ionic compound due to the oppositely charged ions having a lower energy than the original elements.

• As always ENERGY tells us how strongly the ions are attracted to each other. In a solid ionic compound the energy is called lattice energy.

Lattice energy defined• Lattice energy is defined as the

change in energy that takes place when separated gaseous ions are packed together to form an ionic solid.

• It is the energy released when an ionic solid forms from its ions!– Remember the sign is determined from the

system’s P.O.V! Exo = negative: since the energy is leaving the system.

Let’s look at the energy changes involved with the formation of this

ionic solid

• Sublimation – solid to gas phase• Ionization of Li• Dissociation of fluorine• Formation of F- ions• Formation of solid LiF from gaseous ions

2

1Li(s) F (g) LiF(s)

2

This is an energy diagram!

Notice the exothermic reaction…more energy is released than absorbed in the

process

Lattice energy calculations

• This leads us to a modified version of Coulomb’s law:

• Where k is a proportionality constant that depends on the structure of the solid and the electron configurations of the ions and r is the shortest distance between the ions.

1 2Lattice energy=QQ

kr

• Looking at the formula, can you see that the process becomes MORE exothermic as the ionic charges increase AND as the distance between the ions decreases

• Let’s look at how charges affect energy by comparing MgO and NaF energy diagrams.

1 2Lattice energy=QQ

kr

8.6 Partial Ionic Character of Covalent

Bonds• Let’s think of polar covalent bonds and ionic. We understand these bonds to be between atoms that have different electronegativities.

• They either share e- unequally or transfer one or more e-. So how do we tell if it is polar covalent or ionic?

• The REAL answer is there are NO completely ionic compounds. Let’s look.

A complication• The previous slide defies the idea

that we know many of those compounds (that are above 50%) as ionic solid.

• We must consider that the compounds in the table are in the gas phase. And that these results can not be assumed to also apply to the solid phase.

Another complication• Another complication in identifying

ionic compounds is that many contain polyatomic ions.

• Polyatomic ions are held together by covalent bonds

• So calling NH4Cl or Na2SO4 ionic is ambiguous.

What will we call ionic?• So from now on…we define ionic

compound as any compound that conducts an electric current when melted.

8.7 The Covalent Chemical Bond: A

Model• Bonding is a model proposed to explain molecular stability.

• The bond concept is a human invention to provide a method for dividing up the energy evolved when a stable molecule is formed from its component atoms.

Sensible…• It is physically sensible and makes

sense that atoms can form stable groups by sharing e- since shared e- give a lower energy state because they are simultaneously attracted by 2 nuclei.

8.8 Covalent Bond Energies and Chemical

Reactions• From the stepwise decomposition of methane it is determined that the energy to break a C—H bond varies in a nonsystematic way. This also shows that bond strength varies significantly with its environment.

• So bond energy as an average is what is useful to chemists and given in a table format.

Single, double, tripleand I don’t mean baseball

• Single bonds – 1 pair of e-shared.• Double bonds – 2 pair of e- shared.• Triple bonds – 3 pair of e-shared.

• There is a relationship between bond length and # of shared e-. The more e- shared, the shorter the bond.

Bond Energy and Enthalpy

• Using bond energy, the energies of reactions can be approximated.

• Breaking bonds requires energy so the sign is +

• Formation of bonds releases energy so the sign is -

The “formula”• ∆H = sum of the energies required to

break old bonds plus the sum of the energies released in the formation of new bonds.

energy required energy released

∑ is the sum of the terms and D represents the bond energy per mole of bonds. (D is always positive)

(bonds broken)- (bonds formed)H D D

Let’s try with…H2(g) + F2(g) 2HF(g)

We need to break 1) H—H bond 1) F—F bond

We need to form 1) H—F bond

H H F F H FH D D D

432 154 5651 1 2

544

kJ kJ kJmol mol mol

mol mol molKJ

If we compare this with ∆H from standard enthalpy for HF (-271kJ/mol):

∆Ho = 2mol • -271kJ/mol = -542 kJ

So bond energies work well to find ∆H

8.9 The localized Electron Bonding Model

• A “localized electron” assumes that a molecule is composed of atoms that are bound together by sharing pairs of electrons using the atomic orbitals of the bound atoms.

• Lone pairs are the e- that are localized on the atom (or in space)

• Bonding pairs are those e- between the atoms.

Localized electron model when applied

tells us…1. The description of the arrangement of the valence electrons (Lewis structures)

2. A prediction of the molecule’s geometry (VSEPR)

3. A description of the atomic orbital types used to share electrons or hold long pairs. (we’ll talk about this in Chapter 9)

8.10 Lewis Structures• Shows how valence electrons are

arranged among atoms in a molecule.

• Reflects central idea that stability of a compound relates to noble gas electron configuration.

Steps for writing/drawing Lewis

Structures1. Sum the valence electrons. (Add for – ions, subtract for + ions)

2. Use a pair of electrons to form a bond between each pair of bound atoms.

3. Atoms usually have noble gas configurations. Exceptions are Hydrogen (duet rule) and more we’ll talk about later.

8.11 Exceptions to the Octet Rule

• We know that hydrogen is an exception as hydrogen’s outer most orbital is the 1s so when it is full it is 1s2 therefore it follows the duet rule not the octet rule

Another e- poor exception

• Boron – Boron is electron deficient therefore it often only bonds with 6 electrons around it NOT 8. This is indicated often by its violent reactivity.

• Electron poor atoms are typically found in period 2 (this makes sense, because they don’t have 8 e- to begin with). They will NEVER exceed 8 since their 2s and 2p orbitals can’t exceed 8

B – 3e-

3F – 3x7e-

24e-

F B F

F

3 single bonds (3x2) = 69 lone pairs (9x2) = 18

Total = 24

Electron rich exceptions

• There are some occasions where an atom will exceed the octet rule. This is only observed for some elements in period 3 and beyond. Common ones are S and P, but there are more too!

• SF6

S

F

F

F

FF

F

S – 6e-

6F – 42e-

48e-

6 single bonds (6x2) = 1218 lone pairs (18x2) = 36

Total = 48

Should we try?Draw a Lewis structure for each

of the following molecules: 

• H2 • N2 • O2 • F2

• H2O

• NH3

• CO• They get harder as

you go

• CO2

• CH3OH

• *BeCl2• C2H6O

• *PCl5• NO3

-

• *ICl4-

8.12 Resonance• Resonance is when there is more

than one valid Lewis structure for a particular model

• Reality is that the actual structure is an average of the resonance structures.

Here’s one example

O C O

O

O C O

O

OCO

O

2-2-

2-

However,• Resonance is necessary because the

LE model assumes that e- are localized between atoms, however, we know that electrons are really delocalized and they can move around the entire molecule.

• Resonance “compensates” for this defective assumption.

Odd electron molecules• There are very FEW of them!!!• Formed from nonmetals.• The most common example is NO

N – 5e-

O – 6e-

11e-

NO N O

Formal Charge• How do we decide which Lewis

Structure best describes the actual bonding in a molecule.

• One method is to estimate the charge on each atom in the various possible Lewis structures. We will use formal charge to do this.

Things to know:Formal charge is the difference

between the # of valence e- on the free atom and the # of valence e- assigned to the atom in the molecule

To use formal charge we need to know 2 things.

1. The number of valence electrons on the free neutral atom

2. The number of e- belonging to the atom in the molecule.