Covalent Bonding. Covalent Bond: –a bond formed by the sharing of electrons between atoms. (does...

Covalent Bonding

Covalent Bonding• Covalent Bond:

– a bond formed by the sharing of electrons between atoms. (does NOT form charges)

– Made up of nonmetals• Molecule: a neutral group of atoms

joined together by covalent bonds. (Compounds formed with ionic bonds do NOT have molecules)

• Molecular Formula: chemical formula for a molecular compound. It shows how many atoms of each element a molecule contains.

Covalent Bonding• Examples:

ascorbic acid (vitamin C): C6H8O6

C_____, H_____, O_____,

trinitrotoluene (TNT): C7H5N3O6

C_____, H_____, N_____, O_____,

6 8 6

7 5 3 6

Molecular Nomenclature

• Prefix System (binary compounds)

1. Less electronegative atom comes first.

2. Add prefixes to indicate # of atoms. Omit mono- prefix on first element.

3. Change the ending of the second element to -ide.

most

PREFIXmono-di-tri-tetra-penta-hexa-hepta-octa-nona-deca-

NUMBER123456789

10

Molecular Nomenclature

Only use

“mono” on the

second

element.

P2O5 =CO2 =

CO =

N2O =

diphosphorus pentoxide

carbon dioxide

carbon monoxide

dinitrogen monoxide

Naming Covalent Binary Compounds

PCl5N2H4

Cl2O7

IO2

phosphorous pentachloride = dinitrogen tetrahydride

=dichlorine heptaoxide =iodine dioxide =

Lewis Structures

• Electron Dot Diagrams–show valence e- as dots–distribute dots like arrows in an orbital diagram

–4 sides = 1 s-orbital, 3 p-orbitals–EX: oxygen

2s 2pO

X

Lewis Structures

• Octet Rule–Most atoms form bonds in order to obtain 8 valence e-

–Full energy level stability ~ Noble Gases

Ne

Diatomic Molecules• Diatomic Molecule: a molecule

consisting of two atoms.• Diatomic molecules in nature:

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

“HON and the Halogens”

• Structural Formula: represents the covalent bonds by dashes and shows the arrangement of covalently bonded atoms

• Unshared pair or lone pair: a pair of valence electrons that is not shared between atoms.

• Single Covalent Bond: formed by one shared pair of electrons.

• Double Covalent Bond: formed by two pairs of shared electrons.

• Triple Covalent Bond: formed by three pairs of shared electrons.

The halogens form single covalent bonds in their diatomic molecules. Fluorine is one example.

The hydrogen and oxygen atoms attain noble-gas configurations by sharing electrons.

The ammonia molecule has one unshared pair of electrons.

Methane has no unshared pairs of electrons.

Each oxygen atom has two unshared pair of electrons.

This is a double covalent bond.

Oxygen Molecule

Some covalently bonded substances DO NOT form discrete molecules, they form… Network Solids: solids in which all of the

atoms are covalently bonded.

Diamond, a network of covalently bonded carbon atoms

Graphite, a network of covalently bonded carbon atoms

Lewis Structures

Lewis structures are representations of molecules showing all electrons, bonding and nonbonding (lone pairs).

RULES FOR DRAWING LEWIS STRUCTURES (STRUCTURAL FORMULAS OR ELECTRON DOT STRUCTURES)

1. Total the number of valence electrons available in the molecule/polyatomic ion (adjust for charge). 2. Write the symbols for the given formula with some space between them. Put the one with the lowest electronegativity in the center and the others on each of the four sides.

Exception: Hydrogen can NEVER be in the center. If there is a Carbon atom, it is the center atom.

3. Draw one dash between the symbols to represent a bond and fill other octets with lone pairs. 4. Check for full octets, & count the total number of electrons used. The total number of electrons drawn in the structure must equal the number of electrons you counted in step 1.

If too many e-‘s in structure, erase lone pairs and make double and/or triple bonds. If too few e-‘s in structure, expand octet of central atom. Odd number of valence e-‘s in molecule = octet deficient (ie: BF3 & BF2).

5. If the molecule is a polyatomic ion and has an overall charge, put brackets around the structural formula and write the charge outside the brackets in the upper right corner.

Writing Lewis Structures**These are on your reference sheet**

Writing Lewis Structures

1. Total the number of valence electrons available in the molecule/polyatomic ion (adjust for charge)– If it is an anion, add one

electron for each negative charge.

– If it is a cation, subtract one electron for each positive charge.

PCl3

5 + 3(7) = 26

Writing Lewis Structures2. Write the symbols for the given

formula with some space between them. Put the one with the lowest electronegativity in the center and the others on each of the four sides.

• Exception: Hydrogen can NEVER be in the center.

• If there is a Carbon atom, it is the center atom.

3. Draw one dash between the symbols to represent a bond…

Writing Lewis Structures

3. … Fill the octets of all the outer atoms with lone pairs.

Writing Lewis Structures

4. Check for full octets & total number of electrons used.

…If too many e-s used, erase lone pairs and make multiple bonds.

…If too few e-s used, expand central atom octet.

…odd number of valence e-s = octet deficient

Writing Lewis Structures

Writing Lewis Structures

6. If the molecule is a polyatomic ion, put brackets around the structure and write the charge in the upper right corner.

Examples

• Cl2

• O2

• N2

#s of Covalent Bonds

Cl Cl

O O

N N

Single Bond: 2 e- shared

Double Bond: 4 e- shared

Triple Bond: 6 e- shared

Examples

• NH3

• CCl4

• H2S

• SO3

Examples

• ClO4-

Polyatomic Ions and covalent bonding “CHLORATE”

O

Cl O

O

Chlorine has 7 valence e-

Each oxygen has 6 valence e-

When ClO3 comes together they form 3 single covalent bonds

One additional electron completes chlorine with a full valence shell, making this a covalently bonded group with an ionic charge of -1

-1

Examples

• CO32-

O

C O

O

Polyatomic Ions and covalent bonding “CARBONATE”

Carbon has 4 valence e-

-2 Each oxygen has 6 valence e-

When CO3 comes together they form 2 single covalent bonds and 1 double covalent bond

Two additional electrons completes carbon with a full valence shell, making this a covalently bonded group with an ionic charge of -2

Examples

• PO43-

O

O P O

O

Polyatomic Ions and covalent bonding “PHOSPHATE”

Phosphorus has 5 valence e-

-3Each oxygen has 6 valence e-

When PO4 comes together they form 3 single covalent bonds

Three additional electrons completes phosphorus with a full valence shell, making this a covalently bonded group with an ionic charge of -3

VSEPR Theory

Most graphics from: http://wps.prenhall.com/wps/media/objects/602/616516/index.html

VSEPR Theory• VSEPR Theory (Valence-shell

electron-pair repulsion theory): the repulsion between electron pairs causes molecular shapes to adjust so that the valence-electron pairs stay as far apart as possible.

http://wps.prenhall.com/wps/media/objects/602/616516/index.html

2 Charge Clouds

AX2 CO2

3 Charge Clouds

AX2E SO2

AX3 CH2O

Bent or

4 Charge Clouds

AX3E NH3

AX4 CH4

AX2E2 H2OBent or

Quiz

1. What is the geometry around the central atom in each of the following molecular models?

2. How many lone pairs (not shown) are around each central atom?

Above is a ball-and-stick molecular model representation of acetaminophen, the active ingredient in such over-the-counter headache remedies as Tylenol (red = O, gray = C, blue = N, ivory = H):

1. What is the molecular formula of acetaminophen?2. What is the geometry (angles) around each carbon and nitrogen?

(The lines between atoms indicate connections only, not whether the bonds are single, double, or triple.)

3. Indicate the positions of multiple bonds (double & triple) in acetaminophen.

104.5

120

107

120109.5120

120

120

120LP

LP

LP

LPLP

C8H9NO2

Polarity

Bond Polarity

• Most bonds are a blend of ionic and covalent characteristics.

• Difference in electronegativity determines bond type.

E difference: 0.0-0.4

E difference: 0.4-1.7

E difference: >1.7

Bond Polarity

• Electronegativity–Attraction an atom has for a shared pair of

electrons.–higher e-neg atom δ-

–lower e-neg atom δ+

Bond Polarity

• Electronegativity Trend – Increases up and to the right.

Table of Electronegativity

• Nonpolar Covalent Bond–e- are shared equally–symmetrical e- density–usually identical atoms

Bond Polarity

δ+ δ-

Bond Polarity• Polar Covalent Bond

–e- are shared unequally–asymmetrical e- density– results in partial charges (dipole)

• Ionic Bond–e- are transferred

Bond Polarity

• Nonpolar– e- shared equally

• Polar– e- shared unequally

• Ionic– e- transferred

Bond Polarity

E difference: >1.7

E difference: 0.4-1.7

E difference: 0.0-0.4

Bond Polarity

Examples:

• Cl2

• HCl

• NaCl

3.0-3.0=0.0Nonpolar

3.0-2.1=0.9Polar

3.0-0.9=2.1Ionic

+ -

+

Bond Polarity

• Nonpolar Covalent – equally shared e-

• Polar Covalent - partial charges, e- shared unequally

Polar Molecule• One end of the molecule is slightly negative and

the other end is slightly positive• Caused by the presence of a polar bond in the

molecule. (the structure is not symmetrical)• A molecule that has two poles is called a dipolar

molecule, or dipole.

How to Determine Molecular Polarity

Are There Polar Bonds?

Are the polar bondssymmetrical around

the molecule?

Non-PolarMolecule

NO

Check theMolecule Shape

YES

PolarMolecule

NO

Non-PolarMolecule

YES

The water dipole

The ammonia dipole

Self Test

• Is CO2 a covalent or ionic compound?• What is CO2 ’s name?• What is the electronegativity

difference between C and O?• Does CO2 have polar bonds?• Is CO2 a polar molecule overall?

. .O. .

C: :O

CHEMICAL FORMULA/Smallest unitIONIC COVALENT

MolecularFormula

CO2

FormulaUnit

NaCl

COMPOUND

2 elementsmore than 2

elements

TernaryCompound

NaNO3

BinaryCompound

NaCl

ION (charged particle)

1 atom 2 or more atoms

PolyatomicIon

NO3-

MonatomicIon

Na+

Bonding Summary

Ionic Covalent

Octet achieved by:

Transfer of electrons (forming + & - ions)

Sharing electrons

Made of:Metal cation (+) & Nonmetal

anion (-)Nonmetals (above metalloid

line) (no charges)

Characteristics: Brittle Small and very large molecules

Structure:Arranged in alternating

+ & - ions (crystal lattice)Individual molecules

Representative Particle:

Formula Unit: (lowest whole number ratio of

atoms)

Molecule: (group of joined atoms)

Bonding Summary

Ionic Covalent

Physical State: Solid Solid, Liquid, Gas

Melting Point: High Low

Electrical Conductivity:

Yes, when dissolved in water or melted

No

Bonding SummaryMetallic

Octet achieved by:

Valence e- delocalized around metal atoms

Made of:Metal cation (+) & Valence electrons

Characteristics:Maleable, ductile,

lustrous

Structure: “Electron Sea”

Representative Particle: Atom

Metallic

Physical State: Solid (except Hg)

Melting Point: High

Electrical Conductivity: Yes (any form)

Ionic Bonding - Crystal Lattice

Types of Bonds

Covalent Bonding - True Molecules

Types of Bonds

Diatomic Molecule

Metallic Bonding - “Electron Sea”

Types of Bonds

Quiz - answer the following on a sheet of paper1. The following ball-and-stick molecular model is a representation of

thalidomide, a drug that causes birth defects when taken by expectant mothers but is valuable for its use against leprosy. The lines indicate only the connections between atoms, not whether the bonds are single, double, or triple (red = O, gray = C, blue = N, ivory = H):

(a) What is the molecular formula of thalidomide?(b) Indicate the positions of the multiple bonds in thalidomide.(c) What is the geometry around each carbon and each nitrogen?