Co-ordinate Bonds, Intermolecular Forces and

Metallic Bonding

Everything you EVER Everything you EVER wanted to know but were wanted to know but were

afraid to ask!afraid to ask!

Co-ordinate BondingCo-ordinate Bonding

Also referred to as DATIVE bonding.

Occurs when a PAIR of electrons is donated from one atom to another.

Only happens when an empty orbital is present (ie an electron deficient species)

Once the coordinate bond is formed, we treat it exactly the same as a normal covalent bond.

This is how Group 13 atoms get full octets!

Amino-BoraneAmino-Borane

Huh!?!?!?

NH3 – BH3 (don’t worry about the name)

Draw the Lewis Structure for NH3, then for BH3

N has a full octet an a lone pair of electrons.

NH3 is an uncharged molecule.

B only has 6 electrons, it is deficient.

BH3 is an uncharged molecule.

N can “donate” its lone pair to B forming a co-ordinate bond.

Amino-BoraneAmino-Borane

Now we have a new molecule with a covalent bond between N and B.

While there are formal charges on N and B, the overall charge on the molecule is neutral.

The bond formed between N and B is just as strong as any covalent bond.

Co-ordinate BondingCo-ordinate Bonding

Draw the co-ordinate bonds between each of the following pairs:

GaCl3, Cl-

H+, H2O

H+, NH3

BF3, NH3

NH3, GaCl3

PCl3, GaCl3

AlCl3, AlCl3 (to form Al2Cl6)

Intermolecular ForcesIntermolecular Forces

Covalent Bonds are NOT intermolecular forces. They are INTRAMOLECULAR forces.

So what are INTERMOLECULEAR Forces?

The forces of attraction/repulsion that exist between molecules. The ones we need to overcome to change the state of the substance.

Relative Strengths of Bond TypesRelative Strengths of Bond Types

London-Dispersion

Dipole-Dipole

Hydrogen Bonding

Ionic Covalent

Intermolecular Forces

Van der Waals ForcesVan der Waals Forces

Discovered but the Dutch physicist, Van der Waals and named in his honor.

They are weak interactions between molecules are divided into 2 basic types:

1. Dipole – Dipole

When a molecule has a permanent dipole moment (molecule is polar), the negative end of one molecule will sit closer to the positive end of its neighbor.

2. Dispersion or London Forces (weakest of all intermolecular)

Caused by temporary shifts in electron density within a molecule. For instance the 2 e- shared in H2, located on 1 H atom rather than the other.

Dipole – Dipole ForcesDipole – Dipole Forces

+ -

+ -

+ -

+ -

+ -

Ion – Dipole ForcesIon – Dipole Forces

Exist when we dissolve an ionic compound in a polar solvent, like H2O.

The - end of H2O is attracted to a cation (+).

The + end of H2O is attracted to an anion (-).

London ForcesLondon Forces

+ -

2e-

- +

2e-

Caused by the natural vibrations of the electrons in the bond, there is no permanent dipole moment.

The dominant intermolecular force in non-polar substances.

Hydrogen BondingHydrogen Bonding

A specific case of Dipole – Dipole Forces.

Occurs when H is involved in a strongly polar bond with O, N or F.

The H nucleus (just a proton) is attracted to the lone pairs of these highly electronegative atoms.

The result is a network of strong intermolecular interactions between H atoms and available lone pairs.

Occurs extensively in H2O, NH3 and HF.

The primary reason H2O doesn’t behave as almost all other substances in the known universe.

H2O expands when it freezes everything else contracts

Hydrogen BondingHydrogen Bonding

Intermolecular Take-home MessageIntermolecular Take-home Message

The stronger the intermolecular interaction, the stronger molecules are held together.

When molecules are held together more tightly, we see the evidence when we examine physical properties of the substance melting and boiling points increase

Practice QuestionsPractice Questions

For each of the following molecules, state which intermolecular force will be dominant. Then state which (out of each pair) should have the higher melting point.

Molecule 1

Dominant Force

Melting Point

Molecule 2

Dominant Force

Melting Point

H2O H-Bonding 0 ºC H2S Dipole -82 ºC

HCl Dipole -114.2 ºC H2 London -259 ºC

NH3 H-bonding -78 ºC H2O H-bonding O ºC

Br2 London -7 ºC I2 London 114 ºC

MetalsMetals

Metals, they’re an interesting bunch . . .

They play by their own set of rules, so to describe them we need to talk about METALLIC bonding.

First of all what do we know about the properties of metals?

they’re malleable and ductile

they’re electron rich (they give up their electrons to make cations)

they conduct electricity

the free flow of electrons yields many colorful solutions

they’re shiny

they conduct heat

Metallic BondingMetallic Bonding

The “Sea of Electrons” Model

Due to low electronegativities, low effective nuclear charges and large diffuse orbitals, electrons can flow freely from one atom to the next.

Metallic BondingMetallic Bonding

The “Sea of Electrons” Model

Electrons carry electrical current, if electrons can flow freely throughout the metal the metal will conduct electricity.

AlloysAlloys

A mix of 2 or more metals.

2 types: Substitutional and Interstitial

Fe Fe Fe

Fe Fe Fe

C C

Zn Cu Zn

Cu Zn Cu

Brass (Zn and Cu in various proportions)

Typically Zn is put in place of a Cu atom.

Steel (Fe and C in various proportions)

C fills in the “holes” between the Fe atoms.