Chemical Bonding

Chapters 7-8

General Chemistry

Objectives– Explain how atoms combine to form compounds

through both ionic and covalent bonding. – Draw Lewis dot structures for simple molecules.– Relate electronegativity and ionization energy to

the type of bonding an element is likely to undergo.

– Predict the geometry of simple molecules and their polarity (valence shell electron pair repulsion).

Valence Electrons• Valence electrons are the number of electrons in highest occupied

energy level of an atom – The s and p electrons in the outer energy level– Fluorine [He] 2s2 2p5 = 7 valence e-

• The electrons responsible for the chemical properties of atoms are those in the outer energy level

• Core electrons -those in the energy levels below the outer energy level

2s2 2p5

e- Configuration and Valence e-Element

e- config. (use shortcut)

# e- in outer level

Column number from P. Table

Lewis dot charge

Li

Be

B

C

N

O

F

Ne

Lewis Dot (Electron Dot) Diagrams

• Lewis Dot (electron dot) diagrams show valence e- as dots around symbol of element

X

Lewis Dot Diagrams of Selected Elements

Element Valence e- Lewis dot diagram

Sodium 1 Na •

Magnesium

Phosphorus

Chlorine

Octet Rule

• The octet rule: atoms of elements gain, lose or share e- so that each atom has a full outermost energy level

• Want to achieve the e- configuration of a noble gas

• Why named “octet”?

• Exceptions?

Chemical Bonding

• When atoms bond, the valence electrons are redistributed to make the atom more stable

• Ionic bonding: results from the electrical attraction between large numbers of cations and anions

• Covalent bonding: results from the sharing of electrons between two atoms

Ionic Bonding

Remember Ions ?

• Ions: charged atoms

• Cations: positively charged atoms – Metals, like sodium, tend to lose electrons to

create a noble gas configuration (cations)

• Anions: negatively charged atoms– Nonmetals, like chlorine, tend to gain electrons

to create a noble gas configuration (anions)

Ionic Bonds• Formed between metal and nonmetal atoms• Anions and cations are held together by

opposite charges• The bond is formed through the transfer of

electrons • Ionic compounds are called salts• Simplest ratio is called the formula unit

– Example: Na+ will bond with Cl- to make sodium chloride, NaCl

Electronegativity

• Electronegativity: reflects an atom’s ability to attract electrons in a chemical bond

• Metals generally have low electronegativity

• Nonmetals generally have high electronegativity

How Determine if Ionic?

• Ionic bonds form between 2 atoms with difference in electronegativity of 2.0 or greater

Properties of Ionic Compounds

• Conduct electricity in aqueous form– are electrolytes

• High melting and boiling points

• Usually solids at room temperature

• Have crystalline shape

• Example: sodium chloride (table salt)

Lattice Energy

• The strength of an ionic bond compared to another ionic bond is determined by the lattice energy

• Lattice energy is the energy released when one mole of an ionic compound is formed from gaseous ions

• Examples: – NaCl -787.5 kJ/mol (weaker bond)

– MgO -3760 kJ/mol (stronger bond)

Crystalline structure

Ionic Bonding Lewis Dot Diagrams

Na Cl

Ionic Bonding Lewis Dot Diagrams

Na+ Cl-

Ionic Bonding Lewis Dot Diagrams

• All the electrons must be accounted for!

Ca P

Ionic Bonding Lewis Dot Diagrams

Ca P

Ionic Bonding Lewis Dot Diagrams

Ca2+ P

Ionic Bonding Lewis Dot Diagrams

Ca2+ P

Ca

Ionic Bonding Lewis Dot Diagrams

Ca2+ P 3-

Ca

Ionic Bonding Lewis Dot Diagrams

Ca2+ P 3-

Ca P

Ionic Bonding Lewis Dot Diagrams

Ca2+ P 3-

Ca2+ P

Ionic Bonding Lewis Dot Diagrams

Ca2+ P 3-

Ca2+ P

Ca

Ionic Bonding Lewis Dot Diagrams

Ca2+ P 3-

Ca2+ P

Ca

Ionic Bonding Lewis Dot Diagrams

Ca2+ P 3-

Ca2+P

3-

Ca2+

Ionic Bonding Lewis Dot Diagrams

= Ca3P2Formula Unit

Metallic Bonding

Metallic Bonds

• Metallic bonding is the bonding that results from the attraction between metal atoms and the surrounding sea of electrons.– Bond between two metal atoms

+ + + ++ + + +

+ + + +

Sea of Electrons• Metals hold on to their valence electrons very weakly.

• Think of them as positive ions (cations) floating in a sea of electrons

• Electrons are free to move through the solid.

• Metals conduct electricity.

Covalent Bonding

Covalent Bonds

• Two nonmetals share electrons to achieve full octet of electrons

• By sharing, both atoms get to count the electrons toward a noble gas configuration.

• Form molecules - compounds that are bonded covalently

Examples of Molecules

How determine if covalent?

• Covalent bonds form between 2 atoms with difference in electronegativity of less than 2

Properties of Covalent Compounds

• Do not conduct electricity in aqueous solution– Are non-electrolytes

• Relatively low melting and boiling points

• Can be gasses, liquids or solids @ room temp

• Examples: sugar, wax, carbon dioxide

Comparison of MP, BP in Ionic and Covalent Compounds

Bond Energy

• The strength of a covalent bond compared to another covalent bond is determined by the bond energy

• Bond Energy: energy required to break a chemical bond and form neutral, isolated atoms– Stronger covalent bonds have a higher bond

energy

Bond Energy and Bond Length

Bond Length

• Bond Length: the average distance between two bonded atoms

• The longer the bond, the smaller the bond energy (the weaker the bond)

• The shorter the bond, the larger the bond energy (the stronger the bond)

Types of Covalent Bonds

• Single covalent

• Double covalent

• Triple covalent

• Share 2 e- (one pair)

• Share 4 e- (two pairs)

• Share 6 e- (three pairs)

Covalent bonding

• Fluorine has seven valence electrons

F

Covalent bonding• Fluorine has seven valence

electrons• A second F atom also has seven• By sharing electrons…

F F

Covalent bonding• Fluorine has seven valence

electrons• A second atom also has seven• By sharing electrons…

F F

Covalent bonding• Fluorine has seven valence

electrons• A second atom also has seven• By sharing electrons…

F F

Covalent bonding• Fluorine has seven valence

electrons• A second atom also has seven• By sharing electrons…

F F

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

Bonding and Nonbonding Electrons

• Bonding (shared) electrons are involved in a chemical bond

• Nonbonding (unshared or lone pair) electrons are not involved in bonding and belong exclusively to one atom

Nonbonding electronsBonding

electrons

Diatomic Elements

• Seven pure elements that exist as pairs in nature

• Are covalently bonded– H2 N2 O2 F2 Cl2 Br2 I2

• Ways to remember:– BrINClHOF – H, N, O, Halogens

Polarity

Bond Polarity

• Atoms of elements do not always share electrons equally

• Polar covalent bond: unequal sharing of electrons (dif electroneg 0.5 – 1.9)

• Nonpolar covalent bond: equal sharing of electrons (dif electroneg 0.0-0.4)

Bond Polarity• When two different atoms bond

covalently, there is an unequal sharing– the more electronegative atom will have a

stronger attraction and will acquire a slightly negative charge

– called a polar covalent bond or simply polar bond.

Bond Polarity• Refer to Periodic Table values of

Electronegativity• Consider HCl

H = electronegativity of 2.1Cl = electronegativity of 3.0– the bond is polar– the chlorine acquires a slight negative

charge, and the hydrogen a slight positive charge

Bond Polarity• Only partial charges, much less than a

true 1+ or 1- as in ionic bond

• Written as:

H Cl• the positive and minus signs (with the

lower case delta ) denote partial charges.

Bond Polarity• Can also be shown:

• the arrow points to the more electronegative atom.

H Cl

Calculate Polarity of Bond

Difference in Electronegativity

Type of Bond

0.0-0.4 Nonpolar covalent

0.5-1.9 Polar Covalent

2.0 and greater Ionic

Geometry

VSEPR Theory

• Valence Shell Electron Pair Repulsion Theory

• Molecules form in 3-D orientation such that electrons are as far apart as possible

• Allows chemists to predict shapes of simple molecules

Shapes of Molecules

Shapes of Molecules

Predict shape and polarity

• Shape affects polarity of molecule

• Even though atoms may have dif electroneg > 0.5, the shape may cancel out the effects

• Example: CO2

Intermolecular Forces

Intermolecular Forces

• Polar molecules, such as water (H2O) attract other polar molecules.

• The forces of attraction between molecules are known as Intermolecular Forces (IM).– Stronger IM Forces result in higher MP, BP

(solids and liquids)– Weaker IM Force result in lower MP, BP

(liquids, gases)

http://www.langara.bc.ca/biology/mario/Assets/WaterH-bond.jpg

Types of Intermolecular Forces

• Dipole-dipole forces

• Hydrogen bonding

• London Dispersion Forces

Dipole-Dipole

• Dipole-dipole forces– Attractions between polar molecules– Example: BrF

Hydrogen Bonding

• Hydrogen bonding– Is a special type of dipole-dipole attraction– Not really a “bond” but a stronger attraction– Attraction between polar molecules that contain

H bonded to N, O or F

– Example: H2O NH3

London Dispersion Forces

• London Dispersion Forces– Generally only significant IM force in nonpolar

molecules– Attraction between large massed atoms (that

have lots of electrons)– Created by movement of electrons and creation

of instantaneous dipoles

http://www.chem.purdue.edu/gchelp/liquids/disperse.html

References

• http://naio.kcc.hawaii.edu/chemistry/electronegativity.html (electroneg values)

• http://www.oup.co.uk/oxed/children/yoes/pictures/atoms/ (pictures of atoms)

•  http://faculty.gvsu.edu/carlsont/chime/gallery2/gallery2.htm (CO2 molecule)

• Dr. Stephen L. Cotton, Charles Page High School