Bonding, Electronegativity,

& Bond Shape

Forming Chemical Bonds

• According to the Lewis model– an atom may lose or gain enough electrons to

acquire a filled valence shell and become an ion. An ionic bond is the result of the force of attraction between a cation and an anion.

– an atom may share electrons with one or more other atoms to acquire a filled valence shell. A covalent bond is the result of the force of attraction between two atoms that share one or more pairs of electrons.

Types of Chemical Bonding

1. Metal with nonmetal:

electron transfer leads to ionic bonding

2. Nonmetal with nonmetal:

electron sharing leads to covalent bonding

3. Metal with metal:

electron pooling leads to metallic bonding

Typically:

Figure 9.2 The three models of chemical bonding

When drawing Lewis Dot Structures, start with one dot on the top and work your way around.

Kr

Lewis Dot Diagrams

Because valence electrons are so important to the behavior of an atom, it is useful to represent them with symbols. A Lewis Dot Diagram shows dots for each electron in the atom. Each dot represents one valence electron.

Li Be B C N

O F Ne

Ionic compounds form crystals: latticed structures of atoms lined up in predictable ways.

Na Na+ + e-

Cl + e- [ Cl ]-

Na+ + [ Cl ]- Na+[ Cl ]-

Remember, negative and positive charges attract, so Na+ and Cl- stick together, making an ionic compound.

Now chlorine and sodium have 8 valence electrons, fulfilling the octet rule.

Ionic bonding: Al + Cl

[ Cl ]3–[Al]3+

Al + 3Cl [Al]3+[Cl]3–

Cl

Al Cl

Cl

Covalent Bonds in NH3

Bonding pairs

H

H : N : H Lone pair of electrons (unshared pair)

Covalent bondingQ7CCl4 - Covalent

C

Cl

Cl

Cl

Cl

HCl - Covalent

H Cl

MgF2 - Ionic

[ F ]2– [Mg]2+

H2O - Covalent

H O H

NH3 - Covalent

H N H

H

NaCl - Ionic

[ Cl ]– [Na] +

OH– - Covalent

O H

H2 - Covalent

H H

Covalent Bonds

Two nonmetal atoms form a covalent bond because they have less energy after they bonded

H + H H : H = HH = H2

hydrogen molecule

Double Covalent Bond

2 pairs of electrons are shared between 2 atoms

Example O2

O + O O::O

double bond

Triple Covalent Bond

3 pairs of electrons are shared between 2 atoms

Example N2

N + N N:::N

triple bond

Drawing Lewis Structures

1. Determine the number of valence electrons in the molecule

2. Decide on the arrangement of atoms in the molecule

3. Connect the atoms by single bonds

4. Show bonding electrons as a single line; show nonbonding electrons as a pair of Lewis dots

5. In a single bond, atoms share one pair of electrons; in a double bond, they share two pairs, and in a triple bond they share three pairs.

Carbonic acidFormaldehydeAcetyleneEthylene

Hydrogen chlorideMethaneAmmoniaWater

H

H N H C H H ClH

HC C

HC C HH

HC

HHO

H

H2O (8) NH3 (8) CH4 (8) HCl (8)

C2H4 (12) C2H2 (10) CH2O (12) H2CO3 (24)

H

HHO

H

O OC HH

O

Lewis Structures

Electronegativity is a measure of the attraction of an atom for the electrons in a chemical bond. The higher the electronegativity of an atom, the greater its attraction for bonding electrons. Electronegativity is related to ionization energy.

Electrons with low ionization energies have low electronegativities because their nuclei do not exert a strong attractive force on electrons. Elements with high ionization energies have high electronegativities due to the strong pull exerted on electrons by the nucleus.

In a group, the electronegativity decreases as atomic number increases, as a result of increased distance between the valence electron and nucleus (greater atomic radius). An example of an electropositive (i.e., low electronegativity) element is cesium; an example of a highly electronegative element is fluorine.

When two identical atoms form a covalent bond, as in H2 or Cl2, each have equal share of the electron pair. the electron density is the same on both ends of the bond, because the electrons are equally attracted to both nuclei.

But when different atoms bond, as in HCl, one nucleus attracts the electrons from a bond more strongly than the other. The electrons are shared, but unequally. There are partial charges on both sides of the bond, indicated by a lowercase Greek letter delta, . In HCl, the chlorine carries a partially negative charge (-) while the hydrogen carries a partially positive charge (+). This is because the chlorine atom attracts the electrons stronger than the hydrogen.

• To decide whether a bond is covalent or ionic find the difference in electronegativities

< 2.0 covalent

> 2.0 ionic

Try KF, MgS, Cl2

Covalent or Ionic ???

Figure 9.18DEN

3.0

2.0

0.0

Boundary ranges for classifying ionic character of chemical bonds.

Figure 9.19

Percent ionic character of electronegativity difference (DEN).

A bond that carries a partial charge is called a polar covalent bond. Because there are two poles of charge involve, the bond is a dipole. If the entire molecule has a partial charge, it is a polar molecule.

This relative attraction of electrons in a bond is called electronegativity. In HCl, the Cl is more electronegative than the H and attracts the electrons. The electrons spend more time around the Cl than the H.

Looking it up on a periodic table or a similar table, one can see the H has an electronegativity of 2.1 while Cl has one of 2.9. Therefore, the electronegativity difference is 0.8. Electronegativity decreases as you go down a group and increases as you go left to right in a period.

Because the bonding pair in the carbo-fluorine bond is pulled towards the fluorine end of the bond, that end is left rather more negative than it would otherwise be. The carbon end is left rather short of electrons and so becomes slightly positive.

The symbols + and - mean "slightly positive" and "slightly negative". You read + as "delta plus" or "delta positive".

We describe a bond having one end slightly positive and the other end slightly negative as being polar.

The polar nature of the elements sets up an overall charge on the molecule. This overall charge is represented by the symbol and called a dipole moment.

Each dipole moment for each bond faces the - molecule. Then all the dipole moments are “added” and face the general direction of all the other dipole moments.

These 2 molecules have individual dipole moments for each bond, but they face opposite directions, so they cancel each other out for a dipole moment of 0.

3-D Characteristics of Molecules

• Atoms and molecules have 3 dimensions• Shapes of molecules lead to additional properties of

covalent compounds– Polar covalent Bonding

• When electrons are not shared equally between two atoms• Bond that is certain % ionic

– Nonpolar covalent Bonding • Electrons are shared equally• Diatomic atoms

Shapes of Molecules

Number of electron pairs 2

(= negative charge clouds)

Number of bonded atoms 2

Angle 180°

Name of shape linear

Here are some basic compounds that you see everyday. Do you know what they are? As you can see, even though they are made of the same elements (carbon and hydrogen), they have very different shapes. This is due to the way the atoms bond to each other.

In this chart you can see the 6 basic shapes of molecules. Most molecules in the world take one or more of these shapes.

This chart is nice because it shows the number of valence electrons involved in bonding.

Here are a few of the odd molecule shapes you may encounter. The figure shows the non-bonding pairs of electrons. Remember on a Lewis Dot Structure, you may have only a single dot or a double dot. The double dots do not bind to anything. Those double dots are non-bonding pairs.

Na O

Bonding electrons

Non-bonding electrons

As seen in the figure to the left, each shape has a particular bond angle. This angle is determined by how much stress in placed on the bonds. Electrons do not like to be near each other and will push apart as much as possible. Bond angles do not change.

Look at the trigonal planar molecule. The electrons can only push so far before the electrons on the other side begin to push back. These forces keep the bonds in the same place.

VSEPRTheory

Using the VSEPR Model

1. Draw the electron-dot structure2. Identify the central atom3. Count the total number of electron pairs

around central atom4. Predict the electron shape5. Predict the shape of the molecule using the

bonding atoms

Valence Shell Electron-Pair Repulsion Theory or VSEPR

• molecular shape is determined by the repulsions of electron pairs– Electron pairs around the central atom stay as far apart

as possible.• electron pair geometry - based on number of

regions of electron density– Consider non-bonding (lone pairs) as well as bonding

electrons.– Electron pairs in single, double and triple bonds are

treated as single electron clouds.• molecular geometry - based on the electron pair

geometry, this is the shape of the molecule

Figure 10.4

A balloon analogy for the mutual repulsion of electron groups.

VSEPR - electrons (e-) in bonds/orbitals repel each other

4 groups/bonds repulsion yields a tetrahedral shape ~109.5o

3 bonds repulsion yields a trigonal shape ~120o

2 bonds repulsion yields a linear shape 180o

eg: CH4 or NH3, HCO2H, CO2

The steps in determining a molecular shape.

Molecular formula

Lewis structure

Electron-group arrangement

Bond angles

Molecular shape

(AXmEn)

Count all e- groups around central atom (A)

Note lone pairs and double bonds

Count bonding and nonbonding e- groups

separately.

Step 1

Step 2

Step 3

Step 4

SAMPLE PROBLEM 10.7 Predicting Molecular Shapes with Two, Three, or Four Electron Groups

PROBLEM: Draw the molecular shape and predict the bond angles (relative to the ideal bond angles) of (a) PF3 and (b) COCl2.

SOLUTION: (a) For PF3 - there are 26 valence electrons, 1 nonbonding pair

PF F

F

The shape is based upon the tetrahedral arrangement.

The F-P-F bond angles should be <109.50 due to the repulsion of the nonbonding electron pair.

The final shape is trigonal pyramidal.

PF F

F

<109.50

The type of shape is

AX3E

SAMPLE PROBLEM 10.7 Predicting Molecular Shapes with Two, Three, or Four Electron Groups

continued

(b) For COCl2, C has the lowest EN and will be the center atom.

There are 24 valence e-, 3 atoms attached to the center atom.

CCl O

Cl

C does not have an octet; a pair of nonbonding electrons will move in from the O to make a double bond.

The shape for an atom with three atom attachments and no nonbonding pairs on the central atom is trigonal planar.C

Cl

O

Cl The Cl-C-Cl bond angle will be less than 1200 due to the electron density of the C=O.

CCl

O

Cl

124.50

1110

Type AX3

Figure 9.12

The attractive and repulsive forces in covalent bonding.

Bond Energy

• Bond energy or Bond Strength - the energy required to overcome the attraction of covalently bonded atoms.– It is defined as energy required to break bonds in 1

mole of gaseous atoms.– Bond energy depends on the specific elements

involved.– It can vary from molecule to molecule so table values

are averages.

Figure 9.13

Internuclear distance(bond length)

Covalent radius

Internuclear distance(bond length)

Covalent radius

Internuclear distance(bond length)

Covalent radius

Internuclear distance(bond length)

Covalent radius

Bond length and covalent radius.

• Bond length, bond energy, and bond order are closely related– Higher bond order is shorter, and stronger for a given

set of atoms

– With a constant bond order, longer bonds are usually weaker.

SAMPLE PROBLEM 9.2 Comparing Bond Length and Bond Strength

PROBLEM:

PLAN:

SOLUTION:

Using the periodic table, but not Tables 9.2 and 9.3, rank the bonds in each set in order of decreasing bond length and bond strength:

(a) S - F, S - Br, S - Cl (b) C = O, C - O, C O

(a) The bond order is one for all and sulfur is bonded to halogens; bond length should increase and bond strength should decrease with increasing atomic radius. (b) The same two atoms are bonded but the bond order changes; bond length decreases as bond order increases while bond strength increases as bond order increases.

(a) Atomic size increases going down a group.

Bond length: S - Br > S - Cl > S - F

Bond strength: S - F > S - Cl > S - Br

(b) Using bond orders we get

Bond length: C - O > C = O > C O

Bond strength: C O > C = O > C - O

Bond Energies and Hrxn

• In chemical reactions, reactant bonds are broken and new product bonds are formed. The overall energy change of a reaction is the energy change of this process, plus the energy associated with changes of state.

• In gaseous reactions (no state changes)• Hºrxn = Hºbonds broken + Hºbonds formed

Using bond energies to calculate DH0rxn

DH0rxn = DH0

reactant bonds broken + DH0product bonds formed

DH01 = + sum of BE DH0

2 = - sum of BE

DH0rxn

Ent

halp

y, H

Using bond energies to calculate DH0rxn of methane

Enth

alpy

,H

BOND BREAKAGE

4BE(C-H)= +1652kJ 2BE(O2)= + 996kJ

DH0(bond breaking) = +2648kJBOND FORMATION

4[-BE(O-H)]= -1868kJDH0(bond forming) = -3466kJ

DH0rxn= -818kJ

2[-BE(C O)]= -1598kJ

SAMPLE PROBLEM 10.6 Calculating Enthalpy Changes from Bond Energies

SOLUTION:

PROBLEM: Use Table 9.2 (button at right) to calculate DH0rxn for the following

reaction:

CH4(g) + 3Cl2(g) CHCl3(g) + 3HCl(g)

PLAN: Write the Lewis structures for all reactants and products and calculate the number of bonds broken and formed.

H

C H

H

H + Cl Cl3

Cl

C Cl

Cl

H + H Cl3

bonds broken bonds formed

SAMPLE PROBLEM 10.6 Calculating Enthalpy Changes from Bond Energies

continued

bonds broken bonds formed

4 C-H = 4 mol(413 kJ/mol) = 1652 kJ

3 Cl-Cl = 3 mol(243 kJ/mol) = 729 kJ

DH0bonds broken = 2381 kJ

3 C-Cl = 3 mol(-339 kJ/mol) = -1017 kJ

1 C-H = 1 mol(-413 kJ/mol) = -413 kJ

DH0bonds formed = -2711 kJ

3 H-Cl = 3 mol(-427 kJ/mol) = -1281 kJ

DH0reaction = DH0

bonds broken + DH0bonds formed = 2381 kJ + (-2711 kJ) = - 330 kJ

Polarity & Shape

(a) Methane is a nonpolar tetrahedral molecule. (b) Methyl chloride is a polar tetrahedral molecule.

ADD Formal ChargeIn hybridization, add #bonds + #pairs electronsMore clear shape info – arcoMore on resonanceMore on finding polarity by subtracting ENNon-compliance with octet ruleShowing shape with hatch marksAX, AX2, AXE shapes etcParamagnetic vs dimagneticBond order = ½(MO – AO)