Unit 7

Covalent Bonding

Bonding• A metal & a nonmetal transfer

electrons–An ionic bond

• Two metals mix–An alloy (Metallic bond)

• What do two nonmetals do?–Neither one will give away an electron

–So they share their valence electrons–This is a covalent bond

Covalent Bonding• Nonmetals hold on to their valence

electrons• They can’t give away electrons to

bond• Still want to be stable!

– Need noble gas configuration (octet rule)• Get it by sharing valence electrons

with each other.• By sharing, both atoms get to count

the electrons toward noble gas configuration.

Covalent Bonding Fluorine has seven valence electrons A second atom also has seven By sharing electrons… Both end with full orbitals

F F

Covalent bonding Fluorine has seven valence electrons A second atom also has seven By sharing electrons… Both end with full orbitals

F F8 Valence electrons

Covalent bonding Fluorine has seven valence electrons A second atom also has seven By sharing electrons Both end with full orbitals

F F8 Valence electrons

Ways to Illustrate Covalent Bonds• Molecular formula: shows the number

of atoms of each element in a molecule.–Ex. PF3

• Lewis Structures: uses dots to represent bonding between molecular compounds

• Structural Formulas: shows the arrangement of atoms and bonds–Shared electron dots are replaced with a dash

• Models: ball and stick (3-D versions)

Single Covalent Bond• Occurs between nonmetals or a

nonmetal & hydrogen• Sharing of two valence electrons

(1 pair)• Different from an ionic bond –

electrons are SHARED not transferred

An example with dots…•It’s like a jigsaw puzzle•You will be given the formula•You put the pieces together

to make everyone stable or happy –Most atoms need an octet–H & He need a duo–Carbon is often the center

Water

H

O

Each hydrogen has 1 valence electron

and wants 1 moreThe oxygen has 6 valence

electronsand wants 2 moreThey share to make each

other “happy”

Water• Put the pieces together• The first hydrogen is happy• The oxygen still wants one more

H O

Water

• The second hydrogen attaches• Every atom has full energy levels

H OH

Structural formula…

•Replace shared dots with a dash

OH

H

Practice – Dots & Structures

•CH3I

•H2S

•CH2Cl2

•NH3

•C2H6

•SCl2

•AsF3

•SiH4

•CHF3

•

Multiple Bonds

•Sometimes atoms share more than one pair of valence electrons.

•A double bond is when atoms share two pair (4) of electrons.

•A triple bond is when atoms share three pair (6) of electrons.

Carbon dioxide• CO2 - Carbon is

central atom• Carbon has 4 valence

electrons• Wants 4 more• Oxygen has 6 valence

electrons• Wants 2 more

O

C

Carbon dioxide

• Attaching 1 oxygen leaves the oxygen 1 short and the carbon 3 short

OC

Carbon dioxide Attaching the second oxygen

leaves both oxygen 1 short and the carbon 2 short

OCO

Carbon dioxide The only solution is to share more

OCO

Carbon dioxide The only solution is to share more

OCO

Carbon dioxide The only solution is to share more

OCO

Carbon dioxide The only solution is to share more

OCO

Carbon dioxide The only solution is to share more

OCO

Carbon dioxide The only solution is to share more

OCO

Carbon dioxide The only solution is to share more CO2 requires two double bonds Each atom gets to count all the atoms in

the bond

OCO

Carbon dioxide The only solution is to share more CO2 requires two double bonds Each atom gets to count all the atoms in

the bond

OCO8 valence electrons

Carbon dioxide The only solution is to share more CO2 requires two double bonds Each atom gets to count all the atoms in

the bond

OCO8 valence electrons

Carbon dioxide The only solution is to share more CO2 requires two double bonds Each atom gets to count all the atoms in

the bond

OCO

8 valence electrons

Carbon Dioxide

•Replace the shared pairs with dashes

OCO

Practice• O2

• CS2

• CH2O

• N2F2

• NO2

• HCN (triple)

• C2H2 (triple)

Exceptions to the Octet Rule/Patterns of Bonding

1. Some elements with odd number of valence electrons

–BF3

–PCl5

2. Coordinate covalent bonding

Coordinate Covalent Bond• When one atom donates both

electrons in a covalent bond• Carbon monoxide (CO)

OC

Coordinate Covalent Bond When one atom donates both electrons

in a covalent bond. Carbon monoxide (CO)

OC

Coordinate Covalent Bond When one atom donates both electrons

in a covalent bond. Carbon monoxide (CO)

OCOC

Summary of Covalent Bonding• Covalent bonds occur by SHARING

electrons• Occurs between NONMETALS• End product is called a MOLECULE

1. Molecular compound - formed with different elements

2. Diatomic molecules - 2 of the same atom– There are 7 elements that always form

diatomic molecules

–H2 , N2 , O2 , F2 , Cl2 , Br2 , and I2

Diatomic Molecules

Naming Molecular Compounds• Easier than ionic compounds

• No balancing charges

•1 mono-•2 di-•3 tri-•4 tetra-•5 penta-

•6 hexa – •7 hepta –•8 octa –•9 nona –•10 deca –

Naming Molecular Compounds

• 1st element – add the prefix that matches the subscript– Exception – do not add “mono-” if

there is 1 atom– No aa, oo, or ao double vowels

• 2nd element – add the prefix that matches the subscript– Still ends in -ide

Naming• CO2

– 1st element = Carbon; subscript 1•Remember exception•“Carbon”

– 2nd element = oxygen; subscript 2•Prefix for 2 is di-•“Dioxide”

– Full name = carbon dioxide

Practice Naming• S2Cl2

–Disulfur dichloride

• CS2

–Carbon disulfide

• SO3

–Sulfur trioxide

• P4O10

–Tetraphosphorus decoxide

Name Formula• Just look at the prefixes!• Carbon tetrachloride

–1 Carbon, 4 Chlorine atoms

–CCl4

• Iodine heptaflouride• Dinitrogen monoxide• Sulfur dioxide

Common Names

• H2O – dihydrogen monoxide

– Water

• NH3 – Nitrogen trihydride

–Ammonia• CH4 – carbon tetrahydride

• Methane

• HCl – Hydrogen monochloride– Hydrochloric acid

Names to know!

•NH3 - Ammonia

•H2O - Water

•CO – Carbon monoxide

•CO2 – Carbon dioxide

•SO2 – Sulfur dioxide

•CFl4 – Carbon Tetraflouride

Molecular Shapes

• Lewis diagrams & structural formulas are 2-dimensional

• Real molecules are 3-D• If there are 2 atoms, the molecule

has a LINEAR shape (no other options!)–Carbon monoxide (CO)

• If it has more than 2, how do we figure out the shape?

VSEPR Theory

• Valence Shell Electron Pair Repulsion Theory

• Used to predict shape of a molecule

• Negative electrons repel each other and pairs want to be as far apart as possible

Linear• Linear: 2 atoms around central atom, no

unshared pairs on central atom• With three atoms the farthest the two outer

molecules can get apart is 180º.• Will require 2 double bonds or one triple

bond

C OO180º

Trigonal Planar

• Trigonal planar: 3 atoms around central atom, no unshared pair on central atom.

• Angle = 120°

• 4 molecules around a central atom

• All single bonds• Must think in 3-

D!

C HH

H

H

Tetrahedral

Tetrahedral• Tetrahedral: 4

atoms around central atom, no unshared pair on central atom

• A pyramid with a triangular base.

CH HH

H109.5º

So far…SHAPE # SHARED

PAIRS FROM THE CENTRAL ATOM

# UNSHARED PAIRS ON THE CENTRAL ATOM

LINEAR 2 0TRIGONAL PLANAR

3 0

TETRAHEDRAL

4 0

Molecular Shapes

•But what if there are unshared pairs on the central atom?

•They still repel each other…

Bent

OH

H

O HH

<109.5º

• Bent: 2 atoms around central atom, 1 or 2 unshared pair(s) on central atom.

Bent

• Ball and stick model does not show unshared electron pairs

Pyramidal

N HH

H

NH HH

<109.5º

• Pyramidal: 3 atoms around central atom, 1 unshared pair on central atom

Pyrimidial

• Ball and stick model does not show unshared electron pairs

Trigonal Bipyramidal

• Trigonal bipyramidal: 5 atoms around central atom, no unshared pair on central atom

• Angles = 90° and 120°

Adding to the chart…SHAPE # SHARED

PAIRS FROM THE CENTRAL ATOM

# UNSHARED PAIRS ON THE CENTRAL ATOM

BENT 2 1 or 2

PYRAMIDIAL 3 1

TRIGONAL BIPYRAMIDAL

5 0

So how do I determine the shape of a given molecule?

1.Draw the Lewis diagram2.Count the shared and

unshared pairs3.Use the VSEPR Theory to

determine the shape

Which type of bond is it?• Look at which elements are

involved–Metal & nonmetal = ionic bond–2 nonmetals = covalent bond

• Electronegativity – measure of a tendency of an atom to attract a pair of electrons–Influenced by amount of positive charge in the nucleus & electron shielding

Differences in Electronegativity.

• Big difference between values (greater than 1.70) –One atom REALLY wants the electrons and the other…not so much

– Ionic bonding ionic compound• Smaller difference between values

(less than 1.70) –Both have “equal” attraction for the e-

–Covalent bonding molecule

Differences in Electronegativity

•Medium difference–Still a bit of a tug of war over e-–Unequal sharing of electrons–Results in a POLAR covalent bond –Positive and negative poles–Dipole – partially negative on one side, partially positive on the other

Polar Covalent Bond

Differences in Electronegativity

• Very small difference–Share electrons equally

–NONPOLAR covalent bond–No positive and negative poles

Nonpolar Covalent Bond

Polar vs. nonpolar molecules

• Look at polarity of each bond– All nonpolar bonds = nonpolar

molecule (O2)

• Look at the overall shape– Symmetrical polar bonds cancel each

other out so molecule = nonpolar (CO2, CCl4)

– Nonsymmetrical polar bonds = polar molecule (H2O)

Dipole-dipole attraction• Attraction between + part of one

dipole and - part of another dipole• Hydrogen bond - between an

electronegative atom and a hydrogen atom bonded to another electronegative atom –Often involves F, N, or O –Strongest of the intermolecular forces

Hydrogen Bonding

HHO

+ -

+

H HO+-

+

Hydrogen Bonding

Van der Waals – London dispersion force

•Weak intermolecular force caused by negative electrons on one side of a cloud being attracted to a nearby positive nucleus

•Constantly changing

Properties of Molecular Compounds

• Poor conductors of heat & electricity

• Often found as liquids or gases• Weaker attraction between

atoms• Low melting & boiling points

IONIC vs COVALENT

Carbon (Organic) Chemistry

• Carbon plays a dominant role in the chemistry of living things

• Bonding stability– 4 valence electrons– Very unlikely to form ionic bonds– Can form covalent bonds with LOTS of

different elements (especially H & O)• Small molecules link together resulting

in the formation of a large variety of structures often with repeating subunits

Examples of carbon-based compounds

•Simple hydrocarbons•Small carbon molecules with

functional groups•Complex polymers•Biological molecules

Simple Hydrocarbons

• Petrochemicals – Propane, Butane, Octane

Functional groups

• Specific groups of atoms that are responsible for chemical characteristics of a compound

• ALWAYS a close relationship between properties & structure (aspirin, vitamins, insulin)

Complex Polymers & Biological Molecules

• Natural polymers–Proteins, nucleic acids

• Synthetic polymers–Polythene , Polystyrene–Kevlar–Nylon

Common organic molecules

• CH4

• C2H6

• C2H4

• C2H2

• CH3CH2OH

• CH2O

• C6H6

• CH3COOH