Chemical Bonding Chapter 6. Chemical Bonding & Structure Molecular bonding and structure play...

Chemical Bonding

Chapter 6

Chemical Bonding & Structure

Molecular bonding and structure play the central role in determining the course of chemical reactions.

Bonds

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

Bond Energy

- It is the energy required to break a bond.

- It gives us information about the strength of a bonding interaction.

Bond Energies

Bond breaking requires energy (endothermic).

Bond formation releases energy (exothermic).

Chemical Bonds

Chemical Bond

IonicCovalent

Cation Anion Molecule

Ionic Bonds

- Formed from electrostatic attractions of closely packed, oppositely charged ions.

- Formed when an atom that easily loses electrons (metal) reacts with one that has a high electron affinity(nonmetal).

- 2Na(s) + Cl2(g) ----> 2Na+(aq) + 2Cl-

(aq)

Figure 11.8: The structure of lithium fluoride

Figure 11.1: The formation of a bond between two hydrogen atoms

Covalent Bonding

Covalent bonds are formed by sharing electrons between nuclei.

H. + .H ----> H-H

2 hydrogen atoms hydrogen molecule

Types of Covalent Bonds

Polar covalent bond -- covalent bond in which the electrons are not shared equally because one atom attracts them more strongly than the other. A dipole moment exists. HOH, HCl, & CO

Nonpolar covalent bond -- covalent bond in which the electrons are shared equally between both atoms. No dipole moment exists. CO2, CH4, & Cl2

Electronegativity

The ability of an atom in a molecule to attract shared electrons to itself.

As electronegativity increases, the attraction for electrons increases. Fluorine has the highest value at 4.0 and cesium and francium are lowest at 0.7.

08_132

H2.1

Li1.0

Be1.5

Na0.9

Mg1.2

K0.8

Ca1.0

Rb0.8

Sr1.0

Cs0.7

Ba0.9

Fr0.7

Ra0.9

Sc1.3

Y1.2

La-Lu1.0-1.2

Ac1.1

Ti1.5

Zr1.4

Hf1.3

Th1.3

V1.6

Nb1.6

Ta1.5

Pa1.4

Cr1.6

Mo1.8

W1.7

U1.4

Mn1.5

Tc1.9

Re1.9

Np-No1.4-1.3

Fe1.8

Ru2.2

Os2.2

Co1.9

Rh2.2

Ir2.2

Ni1.9

Pd2.2

Pt2.2

Cu1.9

Ag1.9

Au2.4

Zn1.6

Cd1.7

Hg1.9

Ga1.6

In1.7

Tl1.8

Al1.5

B2.0

Ge1.8

Sn1.8

Pb1.9

Si1.8

C2.5

As2.0

Sb1.9

Bi1.9

P2.1

N3.0

Se2.4

Te2.1

Po2.0

S2.5

O3.5

Br2.8

I2.5

At2.2

Cl3.0

F4.0

H2.1

Li1.0

Be1.5

Na0.9

Mg1.2

K0.8

Ca1.0

Rb0.8

Sr1.0

Cs0.7

Ba0.9

Fr0.7

Ra0.9

Sc1.3

Y1.2

La-Lu1.0-1.2

Ac1.1

Ti1.5

Zr1.4

Hf1.3

Th1.3

V1.6

Nb1.6

Ta1.5

Pa1.4

Cr1.6

Mo1.8

W1.7

U1.4

Mn1.5

Tc1.9

Re1.9

Np-No1.4-1.3

Fe1.8

Ru2.2

Os2.2

Co1.9

Rh2.2

Ir2.2

Ni1.9

Pd2.2

Pt2.2

Cu1.9

Ag1.9

Au2.4

Zn1.6

Cd1.7

Hg1.9

Ga1.6

In1.7

Tl1.8

Al1.5

B2.0

Ge1.8

Sn1.8

Pb1.9

Si1.8

C2.5

As2.0

Sb1.9

Bi1.9

P2.1

N3.0

Se2.4

Te2.1

Po2.0

S2.5

O3.5

Br2.8

I2.5

At2.2

Cl3.0

F4.0

Increasing electronegativity

De

cre

asin

g e

lectr

on

eg

ativity

Increasing electronegativity

De

cre

asin

g e

lectr

on

eg

ativity

(a)

(b)

Pauling Electronegativity Values

Electronegativity values for selected elements. See Figure 11.3 on page 334 in Zumdahl.

Percent Ionic Character

%100x

xx (IC)Character Ionic %

A

BA

where xA is the larger electronegativity and xB

is the smaller value. Watch significant figures!!!

Ionic Bond % IC > 50 %Polar Covalent % IC 5 - 50 %Nonpolar Covalent % IC < 5 %

Percent Ionic Character

%100x

xx (IC)Character Ionic %

A

BA

%1003.0

8.00.3 (IC)Character Ionic %

What type of bonding & % ionic character does KCl have?

%1003.0

2.2 (IC)Character Ionic %

73% (IC)Character Ionic %

Ionic

Percent Ionic Character

%100x

xx (IC)Character Ionic %

A

BA

%1003.5

1.25.3 (IC)Character Ionic %

What type of bonding & % ionic character does HOH have?

%1003.5

4.1 (IC)Character Ionic %

40.% (IC)Character Ionic %

Polar covalent

Percent Ionic Character

%100x

xx (IC)Character Ionic %

A

BA

%1003.0

0.30.3 (IC)Character Ionic %

What type of bonding & % ionic character does N2 have?

%1003.0

0 (IC)Character Ionic %

0% (IC)Character Ionic %

Nonpolar covalent

Three Possible Types of Bonds

Nonpolar Covalent(Electrons equally shared.)

Polar Covalent(Electrons shared unequally.)

Ionic(Electrons are transferred.)

Figure 11.2: Probability representations of the electron sharing in HF

Polarity

A molecule, such as HF, that has a center of positive charge and a center of negative charge is said to be polar, or to have a dipole moment.

+

FH

partial positive charge partial negative charge

08_131

F

H

F

F

F

(a)

H F

(b)

H F

H F

H F

H F

The Effect of an electric field on hydrogen fluoride molecules.

08_133

H

O

H

(a)

+

(b)

Dipole Moment for the water molecule.

Polar Water Molecule

The polarity of water allows it to dissolve ionic materials which are essential for life.

The polarity of the water molecule allows water molecules to attract each other strongly (hydrogen bonds). Because of this fact water remains as a liquid at room temperatures and allows the existence of life as we know it.

08_134

HH

N

H

3

(a)

+

(b)

Dipole moment for the ammonia molecule.

08_151

Nonpolar molecule--zero dipole moment.

Achieving Noble Gas Electron Configurations (NGEC)

Two nonmetals react: They share electrons to achieve NGEC.

A nonmetal and a representative group metal react (ionic compound): The valence orbitals of the metal are emptied to achieve NGEC. The valence electron configuration of the nonmetal achieves NGEC.

Noble Gas ConfigurationWhen a Group I, II, or III metal reacts with a

nonmetal to form a binary ionic compound, the nonmetal gains electrons to obtain the configuration of the next noble gas. The metal loses electrons to gain the configuration of the previous noble gas.

Na ----> Na+ + e- configuration of Ne

Cl + e- ----> Cl- configuration of Ar

Noble Gas ConfigurationContinued

When two nonmetals react to form a covalent bond, they share electrons to form noble gas configurations for both.

H. + ----> H--

Hydrogen gains the noble gas configuration of helium, while Chlorine gains the configuration of Argon.

::Cl..

.:Cl

..

..

Anion Size

Anions are always larger than the parent atom because they have added electrons which repel each other. As well, the number of protons is less than the number of electrons so they are not held as tightly.

Cation Size

Cations are always smaller than the parent atom because they have lost an entire electron shell. As well, the number of protons is greater than the number of electrons so the electrons are held tighter.

Relative sizes of some ions and their parent atoms.

Lewis Structure

- 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.

- Developed by G.N. Lewis in 1902.

Lewis Structures

Na. sodium atom [Na]+ sodium ion

sulfur atom [ ] sulfide ion:S..

.

.

2..

..:S:

Lewis Structures

Covalent CompoundsIonic Compounds

::: FF..

..

..

..

In ionic compounds, electrons are transferredand ions are formed. In covalent compounds, electrons are shared to form a molecule. Potassium gains the stability of argon,bromine of krypton, and fluorine of neon.

:BrK1

..

..

1

:

Lone Pairs & Bonding Pairs

:: FF..

..

..

..

Electrons shared between atoms are bonding pairs. Electrons that are notinvolved in bonding are called lone pairs.Each fluorine has three lone pairs and one bonding pair shared between them.

Octet

Neon does not form bonds because it has a full outer shell of electrons--an octet. An octet is four pairs of electrons and represents extra stability for atoms and ions.

Rules for Writing Lewis Structures

• Sum the valence electrons from all the atoms.

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

• Arrange remaining electrons to satisfy the duet rule for hydrogen and the octet rule for the second-row elements.

Lewis Structures

NO+

• 5 e- + 6 e- - 1 e- = 10 e-

• [:NO:]+

• Each atom has an octet and is satisfied.

Single, Double, & Triple Bonds

Single bonds -- one shared pair of electrons.

Double bonds -- two shared pairs of electrons.

Triple bonds -- three shared pairs of electrons.

•Bond Strength = Triple > Double > Single

–For bonds between same atoms, CN > C=N > C—N

–Though Double not 2x the strength of Single and Triple not 3x the strength of Single

•Bond Length = Single > Double > Triple

–For bonds between same atoms, C—N > C=N > CN

Comments About the Octet Rule

- 2nd row elements C, N, O, F observe the octet rule.

- 2nd row elements B and Be often have fewer than 8 electrons around themselves - they are very reactive.

- 3rd row and heavier elements CAN exceed the octet rule using empty valence d orbitals.

- When writing Lewis structures, satisfy octets first, then place electrons around elements having available d orbitals.

Electron Deficient Molecules

Beryllium chloride -- BeCl2 -- is electron deficient with four electrons. It forms a linear molecule.

Boron trifluoride -- BF3 -- is electron deficient with six electrons. It forms a trigonal planar molecule.

See page 351 for the reaction between boron trifluoride and ammonia.

Four Failures of Lewis Structures

Lewis Structures cannot adequately explain:

1. electron-deficient molecules.

2. the paramagnetism of oxygen and other similar substances.

3. odd-electron molecules.

4. resonance.

Odd-Electron Molecules

NO2

• contains 17 electrons.

• cannot satisfy the octet rule.

• a more sophisticated model is needed-the molecular orbital model.

ResonanceOccurs when more than one valid Lewis structure can be written for a particular molecule.

These are resonance structures. The actual structure is an average of the resonance structures called a resonance hybrid.See the resonance structures for the nitrate ion on page 362 in Zumdahl.

ResonanceResonance structures have Lone Pairs and

Multiple Bonds in different positions.

The actual molecule is a combination of all the resonance forms – it does not resonate between the two forms, though we often draw it that way!

•••• •• ••••••••

•• ••O S O O S O•••••• ••••

••••

••••

Stereochemistry

The study of the three-dimensional arrangement (molecular structure) of atoms or groups of atoms within molecules and the properties which follow such arrangement.

VSEPR Model

Valence Shell Electron Pair Repulsion -- The structure around a given atom is determined principally by minimizing electron pair repulsions.

Predicting a VSEPR Structure

1. Draw Lewis structure.

2. Put pairs as far apart as possible.

3. Determine positions of atoms from the way electron pairs are shared.(Parent Geometry)

4. Determine the name of molecular structure from positions of the atoms.(Actual Geometry)

Molecular GeometryParent Geometry is electron pair arrangement about the central atom.

•linear

•trigonal planar

•tetrahedral

Actual Geometry is the arrangement of atoms about the central atom.

•linear

•bent

•trigonal pyramid

08_142

NN

H

H

H

(a) (b)

Lonepair

Lone pair of electrons on the ammonia molecule.

08_143

(a) (b)

H

(c)

Lone pair

Bondingpair

Bondingpair

Lone pair

HH

O

O

H

O

Lone pairs on the water molecule.

VSEPRTwo pairs of electrons are placed 180o apart --

linear arrangement.

Three pairs of electrons are placed 120o apart -- trigonal planar arrangement.

Four pairs of electrons are placed 109.5o apart -- tetrahedral arrangement.

Double bonds and triple bonds count as one effective pair of electrons.

Electron pair arrangement is the parent geometry. Molecular structure is the actual geometry.

Parent & Actual Geometry

When every pair of electrons on the central atom is shared with another atom, the parent and actual geometry are the same.

When one or more pair of electron pairs around a central atom are unshared(lone pairs), the parent and actual geometry are different.

VSEPR Model Summary• Determine the Lewis structure(s) for the molecule.

• For molecules with resonance structures, use any of the structures to predict the molecular structure.

• Sum the electron pairs around the central atom to determine the parent geometry.

• The arrangement of the pairs is determined by minimizing electron-pair repulsions.(Actual Geometry)

VSEPR Model Summary(Continued)

Lone pairs require more space than bonding pairs since they are tightly attracted to only one nucleus. Lone pairs produce slight distortions of bond angles less than 120o.