chapter 9 covalent bonding. 9.1 the covalent bond covalent bond – a chemical bond that results...

Chapter 9

Covalent Bonding

9.1 The Covalent Bond Covalent Bond – A chemical bond that

results from the sharing of valence electrons.– Between nonmetallic elements of similar

electronegativity. A group of atoms united by a covalent bond is

called a molecule. A substance made up of molecules is called a

molecular substance. Molecules may have as few as two atoms.

(CO), (O2)

Polar Covalent Bonds: Unevenly matched, but willing to share.

To describe the composition of a molecular compound we use a molecular formula.

The molecular formula tells you how many atoms are in a single molecule of the compound.– The molecular formula for oxygen is

O2, which tells you that this molecule contains 2 atoms of oxygen.

The molecular formula for sucrose (table sugar) is C12H22O11. This indicates that one sucrose molecule contains 12 carbon atoms, 22 hydrogen atoms and 11 oxygen atoms.

The empirical formula shows the ratio of atoms in the molecule. You get it by dividing all subscripts by the smallest subscript and rounding. C6H12O6 becomes CH2O.

Molecules are often shown using a structural formula.

A structural formula specifies which atoms are bonded to each other in a molecule.

– One type is the Lewis structures.– The Lewis structure uses dots as

before and has a circle drawn around the dots that represent valence electrons that are joined in a covalent bond.

The principle that describes covalent bonding is the Octet rule.

– Atoms within a molecule will share electrons so as to acquire a full set (8) of valence electrons.

If each atom has 7 valence electrons, than they will share one so that they both will have 8. The sharing of that electron is a Single covalent bond.

Single 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

If more than two atoms join in a covalent bond the atoms with the fewest valence electrons will pair up with the atom that has the most up to a total of 8.

If a pair of electrons is not shared they are called an unshared pair.

If two atoms share two pairs of electrons it is called a double covalent bond.

If two atoms share three pairs of electrons it is called a triple covalent bond.

The Lewis structures that you have seen so far use a pair of dots to represent covalent bonds.

Another kind of Lewis structure uses dashes to represent covalent bonds.

– A single dash represents a single covalent bond.

– A double dash represents a double covalent bond

– A triple dash represents a triple covalent bond.

Covalent compounds (molecules) tend to have low melting and boiling points

When an electron pair is shared by the direct overlap of bonding orbitals, a sigma bond results. Single bonds are sigma bonds.

The overlap of parallel orbitals forms a pi bond. Multiple bonds involve both sigma and pi bonds.

Bond length depends on the sizes of the bonded atoms and the number of electron pairs they share.

Bond dissociation energy is the energy needed to break a covalent bond.– Double bonds have larger bond

dissociation energies than single. Triple even larger

Bond length and bond dissociation energy are directly related.

9.2 Naming Molecules– Covalent compounds (Molecules) are named by adding

prefixes to the element names.• Covalent’ (in this context) means both elements are nonmetals.

– The names of molecular compounds will include prefixes that indicate the number of atoms in the molecule. (CO2 would be called Carbon Dioxide, CCl4 is called Carbon Tetrachloride)

– The following prefixes are used:mono-1 di-2tri-3 tetra-4penta-5 hexa-6hepta-7 octa-8nona-9 deca-10

Note: When a prefix ending in ‘o’ or ‘a’ is added to ‘oxide’, the final vowel in the prefix is dropped.

– Remember to use the “ide” suffix. (Cl is chloride not chlorine)

– Exceptions to the rule:• The prefix “mono” is usually not written with the

first word of the compounds name. – CO2 is not called monocarbon dioxide, just carbon

dioxide.

• Some compounds have common names. (H2O is water or dihydrogen monoxide)

Naming Binary Covalent Compounds: Examples

N2S4 dinitrogen tetrasulfide

NI3 nitrogen triiodide

XeF6 xenon hexafluoride

CCl4 carbon tetrachloride

P2O5 diphosphorus pentoxide

SO3 sulfur trioxide

1 mono

2 di

3 tri

4 tetra

5 penta

6 hexa

7 heptaa

8 octa

9 nona

10 deca

* Second element in ‘ide’ from

* Drop –a & -o before ‘oxide’

Naming Acids An acid is a molecular substance that

dissolves in water to produce hydrogen ions (H+).– Acids behave like an ionic compound

because when they dissolve in water, they create cations and anions.

– The cation is always H+ and the anion depends on the particular acid.

• Example: When the acid HCL dissolves in water it forms H+ cations and CL- anions.

• HNO3 acid will form H+ and NO3- ions.

Naming Acids The name of an acid usually results from the name of the

anion(-).– For acids, where the anion ends with the suffix

“ide”(Chloride) the name of the acid begins with the prefix hydro.

– The acids name also includes the root name of the anion(-) and the word acid.

– In addition you need to change the suffix of the anion from “ide” to “ic”.

• By these rules HCL is named hydrochloric acid.

Other acids are named without the prefix “hydro”. The acid HNO3

- derives its name from NO3- which is nitrate. HNO3 is

named Nitric Acid. The following are names and formulas of common acids:

Naming Acids Anion Corresponding Acid

F-, Fluoride HF, hydrofluoric Acid CL-, Chloride HCL, hydrochloric acid Br-, Bromine HBr, hydrobromic acid I-, iodide HI, hyroiodic acid S2

-, sulfide H2S, hydrosulfuric acid NO3

-, nitrate HNO3, nitric acid CO3

-2, carbonate H2CO3, carbonic acid SO4

-2, sulfate H2SO4, sulfuric acid PO4

-3, phoshate H3PO4, Phosphoric acid C2H3O2

- , acetate HC2H3O2, acetic acid

9.3 Molecular Structures Structural formulas (Lewis Structures)

use letter symbols and bonds to show the position of atoms.

Drawing Lewis Structures:

Water

H

O

Each hydrogen has 1 valence electron

and wants 1 more

The oxygen has 6 valence electrons

and wants 2 more

They 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

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 ( I have to tell you)

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 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 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 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 Requires two double bonds Each atom gets to count all the atoms in

the bond

OCO

8 valence electrons

How to draw them To figure out if you need multiple bonds Add up all the valence electrons. Count up the total number of electrons to

make all atoms happy. Subtract. Divide by 2 Tells you how many bonds - draw them. Fill in the rest of the valence electrons to

fill atoms up.

Examples NH3

N - has 5 valence electrons wants 8

H - has 1 valence electrons wants 2

NH3 has 5+3(1) = 8

NH3 wants 8+3(2) = 14

(14-8)/2= 3 bonds 4 atoms with 3 bonds

N

H

N HHH

Examples Draw in the bonds All 8 electrons are accounted for Everything is full

Examples HCN C is central atom N - has 5 valence electrons wants 8 C - has 4 valence electrons wants 8 H - has 1 valence electrons wants 2

HCN has 5+4+1 = 10

HCN wants 8+8+2 = 18

(18-10)/2= 4 bonds 3 atoms with 4 bonds -will require multiple

bonds - not to H

HCN Put in single bonds Need 2 more bonds Must go between C and N

NH C

HCN Put in single bonds Need 2 more bonds Must go between C and N Uses 8 electrons - 2 more to add

NH C

HCN Put in single bonds Need 2 more bonds Must go between C and N Uses 8 electrons - 2 more to add Must go on N to fill octet

NH C

Another way of indicating bonds

Often use a line to indicate a bond Called a structural formula Each line is 2 valence electrons

H HO =H HO

Structural Examples

H C N

C OH

H

C has 8 electrons because each line is 2 electrons

Ditto for N

Ditto for C here Ditto for O

Coordinate Covalent Bond A Coordinate Covalent Bond is 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

Resonance Resonance is a condition when more

than one dot diagram with the same connections is possible.

Resonance structures differ only in the position of the electron pairs, never the atom positions.

9.4 Molecular Shape (VSEPR) Valence Shell Electron Pair Repulsion. Predicts three dimensional geometry of

molecules. The name tells you the theory. Valence shell - outside electrons. Electron Pair repulsion - electron pairs try

to get as far away as possible. Can determine the angles of bonds. And the shape of molecules

VSEPR Based on the number of pairs of

valence electrons both bonded and unbonded.

Unbonded pair are called lone pair.

CH4 - draw the structural formula

VSEPR Single bonds fill

all atoms. There are 4 pairs

of electrons pushing away.

The furthest they can get away is 109.5º.

C HH

H

H

4 atoms bonded Basic shape is

tetrahedral. A pyramid with a

triangular base. Same basic shape

for everything with 4 pairs. CH H

H

H109.5º

3 bonded - 1 lone pair

N HH

H

NH HH

<109.5º

Still basic tetrahedral but you can’t see the electron pair.

Shape is calledtrigonal pyramidal.

2 bonded - 2 lone pair

OH

H

O HH

<109.5º

Still basic tetrahedral but you can’t see the 2 lone pair.

Shape is calledbent.

3 atoms no lone pair

CH

HO

The farthest you can the electron pair apart is 120º.

Shape is flat and called trigonal planar.

Will require 1 double bond

C

H

H O

120º

2 atoms no lone pair With three atoms the farthest they can

get apart is 180º. Shape called linear. Will require 2 double bonds or one triple

bond

C OO180º

Molecular Orbitals The overlap of atomic orbitals from

separate atoms makes molecular orbitals

Each molecular orbital has room for two electrons

Two types of MO

– Sigma ( σ ) between atoms

– Pi ( π ) above and below atoms

Sigma bonding orbitals From s orbitals on separate atoms

+ +

s orbital s orbital

+ ++ +

Sigma bondingmolecular orbital

Sigma bonding orbitals From p orbitals on separate atoms

p orbital p orbital

Sigma bondingmolecular orbital

Pi bonding orbitals P orbitals on separate atoms

Pi bondingmolecular orbital

Sigma and pi bonds All single bonds are sigma bonds A double bond is one sigma and one pi

bond A triple bond is one sigma and two pi

bonds.

Hybridization The mixing of several atomic orbitals to

form the same number of hybrid orbitals.

When an atom approaches another atom to form a bond, the orbitals of its electrons may be changed.

The changing of the orbitals forms what is called hybrid orbitals. This is a cross between all of the orbitals involved (s,p,d,f).

Hybridization When an “s” and “p” orbital come together they

form a linear shaped molecule.

When an “s” and “p2” orbital come together they form a trigonal planer molecule.

When an “s” and “p3” orbital come together they form tetrahedral, pyramidal or bent molecules.

Atoms with higher orbital numbers form longer bonds. Multiple bonds are shorter than single bonds because bonds with more electrons attract the nuclei of the bonding atoms more strongly.

9.5 Electronegativity and Bond Type

Electro negativity

Difference

Bond Type

<.4 Non-Polar Covalent

0.5 – 1.9 Polar Covalent

>2.0 Ionic

Polar Bonds When the atoms in a bond are the

same, the electrons are shared equally. This is a nonpolar covalent bond. When two different atoms are

connected, the electrons may not be shared equally.

This is a polar covalent bond.

Electronegativity A measure of how strongly the atoms

attract electrons in a bond. The bigger the electronegativity

difference the more polar the bond.

0.0 - 0.4 Covalent nonpolar

0.5 - 1.0 Covalent moderately polar

1.0 -2.0 Covalent polar

>2.0 Ionic

How to show a bond is polar Isn’t a whole charge just a partial charge means a partially positive means a partially negative

The Cl pulls harder on the electrons The electrons spend more time near the Cl

H Cl

Polar Molecules Polar molecules have a partially

positive end and a partially negative end Requires two things to be true:

– The molecule must contain polar bonds.This can be determined from differences in electronegativity.

– Symmetry can not cancel out the effects of the polar bonds.

Must determine geometry first.

Polar Molecules Symmetrical shapes are those without

lone pair on central atom – Tetrahedral– Trigonal planar– Linear

Will be nonpolar if all the atoms are the same

Shapes with lone pair on central atom are not symmetrical

Can be polar even with the same atom

Intermolecular Forces They are what make solid and liquid

molecular compounds possible. The weakest are called van der Waal’s

forces - there are two kinds

– Dispersion forces

– Dipole Interactions

Dispersion Force

Depends only on the number of electrons in the molecule

Bigger molecules more electrons More electrons stronger forces

• F2 is a gas

• Br2 is a liquid

• I2 is a solid

Dipole interactions Occur when polar molecules are

attracted to each other. Slightly stronger than dispersion forces. Opposites attract but not completely

hooked like in ionic solids.

Dipole interactions Occur when polar molecules are

attracted to each other. Slightly stronger than dispersion forces. Opposites attract but not completely

hooked like in ionic solids.

H F

H F

Properties of Molecular Compounds

Made of nonmetals Poor or nonconducting as solid, liquid or

aqueous solution Low melting point Two kinds of crystals

– Molecular solids held together by IMF– Network solids- held together by bonds– One big molecule (diamond, graphite)