UNIT 7: BONDING

Why Bond?

Elements bond in order to get a full valence shell

Elements want to have the same number of electrons as the noble gases The Octet Rule: in forming compounds,

atoms tend to achieve the electron configuration of a noble gas

**** only VALENCE electrons are involved in bonding

Types of Bonds

Ionic Covalent Metallic

Review of Metals

Metals lose electrons to form positive ions (cations)

Ionic radius is smaller than atomic radius

Group 1: Alkali metals Lose 1 electron

Group 2: Alkaline Earth Lose 2 electrons

Review of Nonmetals

Gain electrons to form negative ions (anions)

Ionic radius is larger than atomic radius

Group 17: Halogens Gain 1 electron

Group 16: Chalcogens Gain 2 electrons

Ionic Bonds

The force of attraction that holds ions of opposite charge together This attraction is between a metal and a

nonmetal Forms from the TRANSFER of electrons

Electrons are transferred from METAL NONMETAL

Large difference in electronegativity Overall charge of compound is neutral (+

and – cancel out) Ex) NaCl

Animation

http://www.mhhe.com/physsci/chemistry/animations/chang_7e_esp/bom1s2_11.swf

Writing Lewis Dot Structures for Ionic Compounds

Lewis Dot diagrams help us visualize what is happening to valence electrons when a bond forms

Remember- Follow the octet rule!! Ex) Li and F

Steps: 1) Draw all elements and their individual Lewis Dot diagrams. 2) Draw an arrow indicating the transfer of electrons3) Redraw your bonded compound with appropriate charges

More Practice…

Draw the Lewis Dot Diagrams for the following ionic compounds:

BaO

K2S

CaCl2

What if you are not given the formula?

Draw the Lewis Dot diagram for the ionic bonding between sodium and sulfur.

Ionic Compounds

An ionic compound exists as a collection of positively and negatively charged ions arranged in repeating patterns

A formula unit is the lowest whole-number ratio of ions in an ionic compound NaCl (1:1 ratio) , MgCl2 (1:2 ratio)

Properties of Ionic Compounds (salts)

Ionic bonds are the strongest bonds, so all are solids

Hard High melting points and boiling points Soluble in water Cannot conduct electricity as a solid Can conduct electricity as a liquid or in

an aqueous solution http://www.youtube.com/watch?v=NfNIn4R8tg4

They are electrolytes- substances that conduct electricity when dissolved in water

Covalent Bonds

A bond between two nonmetals that involves a sharing of electrons (“tug of war”) Can have EQUAL or UNEQUAL

sharing Small or no difference in

electronegativity Overall charge of a covalent

compound is neutral Covalent compounds are called

molecules Ex) O2, CO, CCl4, H2

Types of Covalent Bonds

There are 3 types of covalent bonds

Covalent Bonds

Non-Polar Covalent

Polar covalen

t

Coordinate Covalent

Non-Polar Covalent Bond

Electrons are shared equally Same atoms- same

electronegativity All diatomics have non-polar bonds! Remember: HOFBrINCl

Two of the same non-metal

Ex) Br2

Polar Covalent Bond

Electrons are shared unequally Different atoms with different

electronegativity values Ex) HF

The more electronegative atom attracts electrons more strongly and gains a slightly (-) charge

The less electronegative atom has a slightly (+) charge

Bond Polarity

• Bond Polarity: refers to a separation of charge in a bond

Ex) HF : partial (+) charge on H and partial (-) charge on F

• This separation of charge is often called a dipole

• ***The greater the difference in electronegativity, the greater the polarity.

Challenge Question

Which of the following covalent compounds has the greatest degree of polarity?

Choice 1) CO

Choice 2) HCl

Choice 3) NO

Choice 4) HBr

Drawing Lewis Structures for Covalent Compounds

Remember! After drawing your diagram, all atoms MUST have 8 valence electrons (except for hydrogen, which should have 2 electrons)

Ex) F2

What if we have more than two non-metals?

Ex) H2O

Ex) NH3

Lewis Structures Continued…

When elements from Group 14 are involved in a covalent bond, they spread their e- out

Carbon tends to form 4 covalent bonds

Ex) CH4

Ex) CCl4

Practice Problem

Draw the Lewis Structure for CH3Br

Can you have more than one bond?

Yes! So far all we have seen are single covalent bonds in which one pair of electrons is shared

Double covalent bond: a bond between two atoms where 2 pairs of e- are shared

Ex) O and O

*** Oxygen tends to form 2 bonds! Triple covalent bond: a bond between two

atoms where 3 pairs of e- are sharedEx) N and N

*** Nitrogen tends to form 3 bonds!

More Practice with Lewis Diagrams..Draw the Lewis diagrams for CO2

Challenge Problem

Draw the Lewis Structure for HCN

Let’s Summarize

C, H, O and N form how many bonds??

Coordinate Covalent Bond

A bond in which both electrons of the shared pair come from the same atom

Ex) NH4+

VSEPR Geometry

Lewis Dot diagrams fail to show the 3-dimensional shapes of molecules

VSEPR Theory (Valence-Shell Electron Pair Repulsion Theory) Repulsion between e- pairs causes

molecular shapes to adjust so that valence e- pairs stay as far apart as possible

http://phet.colorado.edu/en/simulation/molecule-shapes

Molecular Polarity

Molecular polarity is different from bond polarity!!

http://phet.colorado.edu/en/simulation/molecule-polarity In a non-polar molecule, electron

distribution is even (symmetrical) In a polar molecule, electron distribution

is uneven (asymmetrical)

Intermolecular Forces (IMFs)

Only between covalent molecules, never ionic compounds

Weak forces that act between molecules and hold

molecules to each other ****IMFs are not bonds!!!

IMFs occur BETWEEN molecules, bonding occurs WITHIN molecules

IMFs vs. Bonds

Intermolecular Forces of Attraction

• Van der Waals Forces• Dipole interactions• Hydrogen Bonds

London Dispersion Forces

London Dispersion forces: weakest of all molecular interactions; caused by temporary shifts in charge

- Between nonpolar molecules http://antoine.frostburg.edu/chem/senese/101/liquids/faq/h-bonding-vs-london-forces.shtml

- The bigger the atom or molecule the greater the strength of dispersion forces the higher the BP

Think about it…

Which has the strongest London dispersion forces?

F2

Cl2

Br2

I2

Dipole-Dipole Interactions

Dipole interactions: attraction between polar molecules

The positive and negative charges of different molecules attract each other

Ex. HCl http://chemmovies.unl.edu/ChemAnime/DIPOLED/DIPOLED.html

Hydrogen bonds: A special case of dipole-dipole interactions

Hydrogen bonds: intermolecular force between the H of one molecule and a highly electronegative atom of another molecule (must be N, O, or F)

Ex. H2O, NH3

***The high b.p. of water is due to hydrogen bonding

Molecule-Ion Attraction

Partial charges on a molecule are attracted to ions

This is what happens when NaCl dissolves in water

Picture: The hydrogens of water align themselves

towards the anion and the oxygen align themselves towards the cation

http://www.kentchemistry.com/links/bonding/LikeDissolveslike.htm

Molecule-Ion Attraction

Properties of Covalent Compounds Soft Low melting points and boiling

points Cannot conduct electricity in any

phase Generally insoluble in water

Except sugars! (C12H22O11)

Network Solids

A special case of covalent bonding Atoms held together in a very strong covalent

network Ex. Carbon (Diamond) and SiO2 Properties:

Hard High m.p. and b.p. Poor conductors

Holds metals together Electrons are mobile and move from one

atom to another, creating (+) charged metal ions

Charged metal ions are immersed in a “sea of mobile electrons”

Metallic Bonding

Bond Energy

Bonds do not break and form spontaneously- an energy change is required

When a bond is broken, energy is ABSORBED (required) Ex) F2 (g) + ENERGY F (g) + F (g) The greater the # of bonds between

atoms, the more energy you need to break them

When a bond is formed, energy is RELEASED (given off) Ex) F(g) + F(g) F2 (g) + ENERGY