JD N S J TN CB15Metallic Bonding

Definition: Electrostatic force of attraction between metal cations and sea of delocalized electrons in a metal.

Factors affecting strength of metallic bond

Number of valence electrons contributed per metal atomThe larger the number, the stronger. (more electrons=greater attraction with cations)

Charge, radius and charge densityHigher the charge and smaller the radius, the higher the charge density, stronger the bond

Physical Properties

Melting and Boiling Points HighStrong electrostatic forces of attraction between metal cations and sea of delocalized electrons

Good electrical/heat conductivityHas sea of delocalized electrons and ions acting as mobile charge carriers. Delocalised electrons can also pick up heat energy to vibrate faster.

Shiny SurfaceWhen photon strikes surface, oscillating electric field causes electrons to oscillate. Photon bounces off without loss in momentum.

HardDepends on metallic bond strength

Malleable and DuctileLayers of atoms can slide over one another into new positions

Allow Combinations

BronzeCopper and Tin

BrassCopper and Zinc

SteelIron and Carbon

Stainless SteelIron and Chromium

Ionic Bonding

Definition: Electrostatic force of attraction between positively charged cations and negatively charged anions.

Factors affecting strength of metallic bond

Charges on cations and anionsThe larger the number, the stronger.

RadiusSmaller the radius, the higher the charge density.

LATTICE ENERGY=(Cc + Ca) / (Rc + Ra)

Physical Properties

Melting and Boiling Points HighStrong electrostatic forces of attraction between cations and anions in a giant ionic lattice

Good electrical when dissolved in waterDissociate into mobile ions and act as mobile charge carriers to conduct electricity

Hard and rigid but brittleStrong force causes like charges to come to next to each other-repel, shatter

Usually Soluble in WaterIon forms ion-dipole interactions with water, releasing energy in the process, and if this energy is greater than the ionic bonds, the ions become detached from crystals surface.

Covalent Bond

Definition: Electrostatic force of attraction between nucleus of each of 2 bonded atoms and shared pair of electrons.

Sigma/Pi Bonds

Sigma Bond

HEAD-ON OVERLAP between (S-S/S-P/P-P)

Pi Bond

Side-Side Overlap

(Single bond 1 sigma, double bond 1 sigma 1 pi, triple bond 1 sigma 2 pi)

VSEPR

Linear180

Bent118/104.5

Trigonal Planar120

Tetrahedral109.5

Trigon pyramidal107

Octahedral 90

Square Planar90

General Steps for Drawing VSEPR1. Determine number of pairs of electrons, be it bond or not2. Determine how many is lone pair3. Find Structure for total pairs of electrons (e.g four pairs means tetrahedral)4. If there are lone pair, take away one pair and push the rest closer together to form your shape.5. Visualize!

Hybridization

Consider this methane molecule. It has a central C atom.

Carbon forms 4 bonds in CH4, and as you can see its orbital diagram, it only has 2 lone electrons in the p orbital, so it cannot form 4 bonds with the 4 hydrogen!

AND SO IT HAS TO HYBRIDIZE, WHICH MEANS

This happens, one of the electrons from 2s will fill up the last 2p orbital.

And since hybridization means the s and p orbitals join together to form hybridized orbitals:

Why is it SP3 this time round?

Because 1 S orbital combines with 3 P orbitals

They will look like this. As they have 75% P orbital shape and 25% S orbital shape. This is probably the simplest way to visualize it.

And then we ask: Is there an easier and faster way to determine SP/SP2/SP3. Definitely! The previous way was just to explain to you the concept.

Remember this: 1. All SIGMA BONDS are made from hybridized orbitals. 2. All lone pairs are made from hybridized orbitals. 3. All p bonds are made from LEFT-OVER P orbitals (if you understood the concept above)

Lets say we have BF3. How do we determine hybridization of B?

S

P2

It has no lone pairs and 3 sigma bonds. Therefore it is SP2 hybridized.

H2O has 2 lone pairs and 2 bonds for the O central atom. Therefore it is (S+P+2+3) = SP3 hybridized.

ELECTRONEGATIVITY

Weve studied about IONIC/COVALENT bonds (complete electron transfer and equal electron sharing), but have we thought about atoms that are in the middle with both ionic and covalent character? This is where electronegativity comes in.

Definition: Ability of an atom to attract the electrons in a covalent bond to itself.

Increases across the periodEffective nuclear charge increases across the period, stronger electrostatic forces of attraction between nucleus and outermost electron, much easier to attract electrons.

Decreases down the group.Effective nuclear charge decreases due to increasing shielding effect due to increasing quantum shells.

AND SO, we have unequal sharing of electrons, which brings about polar bonds, in every tug of war game someone will win!

IONIC BONDS WITH COVALENT CHARACTER (SUB TOPIC)

We KNOW some cations with very significant charges and very small radiuses (high charge density)These will attract the electron cloud of the anion.

We say they have high polarizing power.

Some anions are very large too, so the electron cloud is very far from the nucleus, and is less attracted.This will allow the cloud to be easily pulled away by the cation.

We say they have high polarisability!

This is WHY some metal non-metal compounds are considered covalent! (e.g BeCl2 AND AlCl3)

And this is also why you might have heard that electronegativity determines whether a bond is PREDOMINANTLY IONIC or PREDOMINANTLY COVALENT.

COVALENT BONDS: formed when atoms of metals and non-metals respectively have similar electronegativity (and so will share electrons EQUALLY)

IONIC BONDS: formed when atoms of metals and non-metals have great difference in electronegativity (one will pull ALL the electrons towards itself!)NET DIPOLE MOMENT

As we can see in C2CL4. It has polar C-CL bonds, but the overall molecule has no net dipole MOMENT. This is because its bond dipole moments cancel out. Use vectors.

This is why TETRAHEDRAL molecules ARE ALWAYS NON POLAR!

PHYSICAL PROPERTIES OF COVALENT COMPOUNDS

Low melting and Boiling PointsLess energy needed to overcome weaker dispersion/dipole attraction/hydrogen bonding forces.

Electrical conductivityNo mobile charges, unable to conduct electricity as all electrons held in covalent bonds and cannot move.

SolubilitySomething is soluble if energy released from solute-solvent interactions higher than energy needed to overcome solute-solute interactions and solvent-solvent interactions.

Polar solutes will dissolve in polar solvents, same applies for non-polar solutes. Water is polar.

Giant Covalent Structures (Learnt in Sec 4 Diamond, SIO2, Graphite)Strong Covalent bonds. Graphite can conduct electricity since its carbon is only bonded to 3 others, leaving 1 delocalised electron to act as mobile charge carriers. Its layers can also slide over one another easily.

INTERMOLECULAR FORCES

*Differentiate this between your ionic/covalent/metallic bonding! This is the force that holds molecules together, not individual atoms or ions! This is also the determinant for your melting/boiling points.

Dispersion Forces-Instantaneous dipole-induced dipole (between non-polar molecules, remember that electrons are never stationary and always moving, so at some random point of time more may be on one side, causing an instantaneous dipole which induces a dipole in another non polar molecule) (This is short lived, since it works based on randomness.)

*POLAR MOLECULES experience instantaneous dipole-induced dipoles too, in addition to their permanent dipole-permanent dipole attractions!

Factors affecting StrengthNo. of electrons (more important)-Larger molecules with more electrons have larger electron clouds, being more polarizable, forming more instantaneous and induced dipoles, resulting in stronger dispersion forces.

Surface area of contact btw adjacent molecules (same number of electrons)-Larger surface area allows more points of contact over which electron clouds can be distorted, so induced dipoles can form easily.*Here you can see both have same number of electrons but n-pentane has higher boiling point as it has a larger surface area.

Permanent dipole-permanent dipole attractions-Between POLAR molecules. Positive dipole of one molecule attracts negative dipole of another.

Hydrogen Bonding-Stronger type of permanent dipole-permanent dipole attraction

That is why it requires TWO polar molecules too!

-Requires H bonded to FON with another FON that has a lone pair

-Why H20, NH3, HF have different boiling points then (consider electronegativity and limiting factor of number of lone pairs and H bonded to FON)

-Water to have high surface tension- Higher density in solid than liquid. Ice less dense than water!

VERY IMPORTANT THINGS YOU HAVE TO TAKE NOTE!

1. Smaller atoms form stronger bonds. Overlap between orbitals more effective and orbitals closer to nucleus.

2. Polar bond stronger and shorter than non-polar bond. Electron more attracted to a nucleus, requires more energy to break, bond length shorter.

3. Multiple bonds stronger than single bonds. (More shared electrons, more attraction between nucleus and electrons)