Bonding Considerations The following PowerPoint presentation is a self directed study of additional…

The following PowerPoint presentation is a self directed study of additional bonding considerations. It includes information above and beyond the ionic and covalent bondingthat you have learned in class.

You are expected to review and study this PowerPoint on you own time.When you feel you are ready, ask your teacher for a handout that will testyour knowledge of these additional aspects of bonding nature.

Let’s ReviewIONIC BONDSIONIC BONDS exist when one atom transfers

electron/s to another atom. Both atoms have attained a Noble Gas configuration.

COVALENT BONDS exist when two atoms share a pair of electrons between them in what is known as a SIGMA ()bond. It is also possible for atoms to share more than one pair of electrons. These multiple bonds are known as PI ()bonds. The atoms involved in bonding have attained a Noble Gas configuration.

VSEPR Theory ExpandedConsider the tetrahedral sp3 hybrid. It

produces an atom with four bonding orbitals each separated by 109.5o.

However, this angle can be altered, in particular, by unshared electron pairs. Let’s consider two common examples of this phenomena.

WaterThe bonding angle between the two Hydrogens in H2O with

its two unshared electron pair lobes on the Oxygen is 104.5o. WHY?

The two unshared pairs, are not confined (restricted) between the nuclei of two atoms.

Therefore, they are able to expand to a greater degree and “squeeze” together the bonds of the two Hydrogen atoms.

AmmoniaThe bonding angle between the three Hydrogens in NH3

with its unshared electron pair lobe is 107o. WHY? The unshared pair on the Nitrogen is not confined

(restricted) between the nuclei of two atoms.Therefore, it is able to expand to a greater degree and

“squeeze” together the bonds of the three Hydrogen atoms.

How can I see these angle changes for myself?

1) Using your Lewis structure directions, make a diagram of CH4 and NH3 and H2O.

2) Note the unshared pairs of electrons on the ammonia and water molecules.

3) Using the ball and stick sp3 models make models of methane, ammonia, and water.

4) Note that in the case of the ammonia and water there will be one and two holes respectively that aren’t filled. These are the unshared pairs. Fill them with gray sticks to represent the unshared electron pairs.

I’m having trouble visualizing these molecules. What can I do?Methane: Note four sigma bonds with the H

atoms which complete the octet of the Carbon.

I’m having trouble visualizing these molecules. What can I do?Ammonia: Note the three sigma bonds with

the H atoms. The lobe of unshared electrons are from the valence electrons of the N atom itself.

I’m having trouble visualizing these molecules. What can I do?Water: Note the two sigma bonds with the H

atoms. The two lobes of unshared electrons are from the valence electrons of the O atom itself.

Note the two lobes of unshared electrons on the lower right of the water molecule.

What have you learned so far?QUESTIONS:1) What are all four angles for a tetrahedron?2) What is the angle between the two H’s in a water molecule?3) What are the angles between three H’s in an ammonia

molecule?4) Why are the angles in water and ammonia less than a

tetrahedron?ANSWERS:1) 109.5o

2) 104.5o

3) 107o

4) The lobes of the unshared pairs expand, take up more space, and push the bonded atoms closer together.

Bond Strength ~ INTRAMOLECULAR BONDS

COVALENT ~ Bonds within a single molecule

IONIC ~ Bonds within a single crystal

Bond strength is determined by BOND ENERGY, i.e. the amount of energy to break the bonds in Kilojoules per mole of bonds

RANK ~ Strongest to Weakest

1) Network Covalent Bonds

2) Ionic Bonds3) Metallic Bonds4) Polar Covalent Bonds5) Non-polar Covalent

Bonds

NETWORK COVALENT SOLIDS

Examples: C (diamond), SiO2 (silica sand)Structural Particles: AtomsElectronegativity Difference: ZeroForces between Particles: Non-polar covalent

bondsProperties: Hard, very high-melting solids;

nonconductors; insoluble in common solvents

Can I see a model of a diamond?Here it is: C-C

bonding!Here’s SiO2 for you

too! SiO2 is sand and

quartz

IONIC BONDSExamples: NaCl (table salt), CaCO3 (calcite)Structural Particles: ions (cation and anion)Forces between Particles: Ionic bondsElectronegativity Difference: >1.7Properties: High melting points; conductors

in the molten state or water solution; usually soluble in water; insoluble in organic solvents

Ionic Crystals can have many different shapes.Table Salt

Copper sulfate

Triphylite

METALLIC BONDSExamples: (sodium), Fe (iron), Au (gold)Structural Particles: cations and mobile

electronsForces between Particles: Metallic bondsProperties: Variable melting points; conductors

in solid state; insoluble in common solvents

LEARN MORE ABOUT THE NATURE OF METALLIC BONDS BY CLICKING THIS LINK: www.ausetute.com.au/metallic.html

A typical metal

POLAR COVALENT BONDSExamples: NH3 (ammonia), HCl

(hydrochloric acid)Structural Particles: Polar moleculesForces between Particles: Polar covalent

bondsElectronegativity Differences: 0.2 – 1.7Properties: Generally higher melting points

and boiling points than non-polar molecules; more likely to be water- soluble

NON-POLAR COVALENT BONDSExamples: H2 (hydrogen gas), CCl4 (carbon

tetrachloride)Structural Particles: Non-polar moleculesForces between Particles: Non-polar covalent

bondsElectronegativity Differences: 0 – 0.2Properties: Low melting and boiling points; often

gas or liquid at 25oC; insoluble in water; soluble in inorganic solvents

What have you learned so far?QUESTIONS:1) Intramolecular bonds are between what types of particles?2)What is bond energy? What unit is used?3) Rank the bond strength from highest to lowest.4) How are metallic bonds different from the other bonds?ANSWERS:

1) Within a single molecule (covalent compounds) & within a single crystal (ionic bonds)2) It is the energy to break a mole of bonds and the unit is kJ/mol.3) Coordinate covalent, ionic, metallic, polar covalent, non-polar covalent4) They involve a moving “sea of electrons.”

Bond Strength ~ INTERMOLECULAR BONDS

These are bonds between molecules (covalent compounds).

Technically they are not bonds in the normal sense. Rather, they are attractive interactions.

The bond strength is measured in Kilojoules per mole of bonds.

Together these attractions are known collectively as van der Waals forces.

RANK ~ Strongest to Weakest

1) Hydrogen bonds2) Dipole – dipole

interactions3) Dipole – induced dipole

interactions4) Dispersion (London)

forces

HYDROGEN BONDSA force exerted between an H atom bonded to an F,

O or N atom on one molecule and an unshared electron pair on the F, O or N atom on another molecule.

The H on the molecule behaves almost like a bare proton because of the high electronegativities of the F (4.0), O (3.5) and N (3.0).

The small size of the H atom allows the unshared pair of the F, O or N to approach very closely. NOTE: This only happens with these three non-metals with their small atomic radii.

H bonding creates relatively high melting and boiling points compared to the low molar masses.

Water is affected by hydrogen bondingThe high surface tension of water is due to H

bonds.NOTE: H bonding creates higher melting and

boiling pointsWater, H2O,\; H bonds:

b.p. 100o C molar mass 18 g/mol

Methane, CH4; no H bonds:b.p. -161.6oC molar mass 16 g/mol

DIPOLE-DIPOLE INTERACTIONSPolar molecules experience an asymmetrical

electronegativity difference higher than 0.2 across the molecule.

Such a molecule is called a dipole.The dipoles line up as close as possible, positive

end to negative endThere is an attractive force between adjacent

molecules. This is known as the dipole moment. Mr. Congdon’s Joke: What did the two dipoles say to each other? You got a moment?

DIPOLE-INDUCED DIPOLE INTERACTIONSA permanent dipole molecule such as

Hydrogen fluoride (H-F) can “induce” (create) a temporary dipole moment in an adjoining non-polar molecule

The overall electron cloud of the molecule (or part of the molecule) will shift to create (+) and (-) poles

Note that these interactions are temporary and the non-polar molecule will shift its electron cloud position back to normal once the permanent dipole is no longer close

Hydrogen flouride

DISPERSION (London) FORCESThese forces involve attractions between temporary

or induced dipoles in adjacent molecules.At a given instant, the electron cloud around a non-

polar molecule may shift from one side of the molecule to the other, thus inducing a dipole.

The temporary dipole induces a similar dipole in an adjacent non-polar molecule.

ALL molecules have dispersion forces. The strength of the forces depends on…..the number of electrons that make up the moleculeThe ease with which the electrons are dispersed to form

temporary dipoles

London forces are responsible for holding much of your body together because of the many molecules involved. Have you every split your lip or your knee bumping into

something? OUCH! You exerted enough bond energy to break the London forces in

your lip or in your knee. Remember London forces are the weakest of the intermolecular

forces.

Once you have reviewed this PowerPoint, see your instructor for an evaluative problem set.

1) You will be asked to fill out several questions related to this PowerPoint.

2) You may refer back to the PowerPoint at any time.

3) The evaluative problem set will also give you book references.

4) You will also be asked to evaluate this PowerPoint and offer suggestions.