Bonding

Metallic

Ionic

Covalent

The attraction between two oppositely charged entities

Metallic•Formed between

Metals•Caused by

The metals’ outer shell electrons being delocalised within the lattice of the resultant metal ions•Actual Bond

The attraction between the +ve ions and the –ve electrons•Resultant structure:

A lattice of +ve ions (usually a cube) in a ‘sea’ of mobile delocalised electrons

Properties of Metals•High mp

Strong attraction between the +ve ion and –ve electron.

•Conductivity

Delocalised mobile electrons are able to move towards a +ve plate.

•Malleable

The layers of ions are able to move over each other.

Structure of metals• Giant Structures – HIGH MP / BP• Strong electrostatic forces of attraction ( metallic bond) between every

particle in the structure.

• How does the metallic bond get stronger?Increase the charge on the ion (stronger attraction between ion and

electron, so stronger metallic bond)

How Does MP/BP vary from Na Mg ?

• How do we increase the conductivity of metals?Increase the number of free electrons

How Does conductivity vary from Na Mg ?

Increases as Charge on ion increases from +1 +2

Increases as number of delocalised electrons per mole of lattice increases from 1 2

Ionic•Formed betweenMetals and Non-metals•Caused byThe metals’ outer shell electrons being transferred to the outer shell of the non-metal so that both achieve a full outer shell. The metal atom becomes a +ve ion and the non-metal atom a –ve ion•Actual BondThe attraction between the +ve ions and the –ve ions•Resultant structure: A lattice of +ve ions (usually a cube) surrounded by –ve ions

FLi

LiF

1

2

4 35

6

Represents the electrostatic attraction between opposite charged ions

Properties of ionic compounds•High mp

Strong attraction between the +ve ion and –ve ion.

•Conductivity in solids very poor

Attraction between the +ve and –ve ions hold the ions in a fixed position in the lattice.

•Conductivity in aqueous solutions or liquids very high

The ions become mobile so +ve ions move towards –ve plates and –ve ions move towards +ve plates.

Structure of Ionic compounds• Giant Structures – HIGH MP / BP• Strong electrostatic forces of attraction ( ionic bond)

between every particle in the structure.

• How does the ionic bond get stronger?

Increase the charge on the ions

(so that the attraction between the ions increase,

stronger ionic bond)

How Does MP/BP vary from NaCl Na2O MgO?

Increases as ions charge increase from +1/-1 +1/-2 +2/-2

Polarised Ionic•Small Highly charged cation

•(Nucleus of cation is not shielded greatly)

•Large anion

•Outer shell of anion is shielded from its’ nucleus

RESULT:

Outer shell of anion is attracted towards the nucleus of the cation

Distorts (polarises) the shape of the anion

Polarised ionic

9+3+

+ -

3+

+

-

53+53+

•Formed between

Non-Metals.•Caused by

The un-paired non-metals’ outer shell electrons are shared to form a electron pair between the two atoms.•Actual Bond

The attraction between the shared electron pair and the two nuclei.•Resultant structure:

Small discrete molecules or large giant structures

Covalent

What is really happening to the electrons during covalent bonding?

• GCSE = Shells overlap• A level Electrons are in (atomic) orbitals • In bonding the (atomic) orbitals overlap

QUESTION: WHAT DO THE ORBITALS LOOK LIKE?

S =

3 p = The 3 p orbitals are all at right angles to each other

Molecular orbitals

• Found in a covalent bond • The region of space in which the shared electron

pair in a covalent bond is located• Formed by the overlapping of the ATOMIC

ORBITALS• Two types

IN ALL SINGLE BONDS THE ELECTRON PAIRS ARE LOCATED IN SIGMA MO

IN ALL DOUBLE BONDS 1 ELECTRON PAIR IS LOCTAED IN A SIGMA MO AND THE OTHER ELECTRON PAIR IS LOCATED IN A PI MO

s and sSigma (σ) p and p end onSigma (σ) p and p side onPi (π)

Sigma MO Pi MOSigma MO

Properties of covalent compoundsSmall discrete Molecules

•Low mpWeak attraction between the small molecules

•Poor Conductivity No charged particles free to move

SIMPLE COVALENT (MOLECULAR) STRUCTURES

Properties of covalent compoundsGiant structures•High mp

Strong covalent bond between every atom in structure

•Low Conductivity

No charged particles free to move

EXCEPTION

Graphite – Delocalised electrons within each layer enables conduction.

Strange Compounds

• GCSE idea: After sharing electrons, atoms posses a full outer shell

• Dot cross diagrams for NaCl, CaO, CO2, HCl, SO2, SO3, BH3

• Note the problems with last three!

NaCl

CaO

CO2,

HCl

SO2

SO3

NOTE S has 10 electrons in outer shell

NOTE S has 12 electrons in outer shell

BH3

NOTE B has 6 electrons in outer shell

What causes Bonding to happen?Electronegativity

The ability of an atom to attract electrons towards itself in a covalent bond

AFFECTED BY?

a. No of shells (and hence the shielding of the nuclei)

b. No of protons in the nucleus

Electronegativity and Periodic Table• Across a PeriodNo of shells the sameShielding the sameNo of Protons increasesNuclear pull on electrons increasesElectronegativity increases

• Down a GroupNo of shells increasesShielding increasesNuclear pull on electrons decreasesElectronegativity decreases

MOST ELECTRONEGATIVE ATOM = FLEAST ELECTRONEGATIVE ATOM = CsELECTRONEGATIVITY OF C < ELECTRONEGATIVITY OF HALOGENSELECTRONEGATIVITY OF C = ELECTRONEGATIVITY OF HYDROGEN

Consider a Covalent bond X-Y

Elec X = Elec Y

Elec X > Elec Y

Elec X >>> Elec Y

ElecX >>>>>>>>>>Elec Y

Hence

Bond between a metal (low elec) and a non-metal (high elec) usually IONIC

Bond between 2 non-metals (similar elec) usually Covalent

Polarised covalent or ionic bonds ensue if somewhere in between

IONIC BOND= No overlap of shells

COVALENT Bond = Overlap of shells

Boron Compounds

• Boron is in group 3

• Expect ionic compounds (lose 3 electrons to achieve a full outer shell)

• Electronic configuration of B+3 = 1s2

• Polarisation so great: Covalent bonds ensue (even with BF3)

NB: B-F bond electron pair nearer to F atom once covalent

Aluminium compounds

• Al+3

• Electronic configuration = 1s22s22p6

• Most compounds form a degree of polarised ionic (except for AlF3)

• All aluminium tri-halides (except for AlF3) are covalent molecules

Covalent Bonding

1.Shared electron pair(s) of electrons between 2 nuclei

2.Pair usually results from 1 unpaired electron from one atom, pairing up with an unpaired electron from another atom. Doesn’t have to be the case

3.Two types of electron pairs are seen in molecules.

Square = Bonding Electron Pair Circle= Non-Bonding Electron Pair or LONE pair

Types of electron pairs in a molecule

Dative Covalent Bonding•2 H2 + O2 2 H2O•H+ + OH- H2O

One of the bonds in the water formed in the lower example, must involve a shared electron pair in which BOTH electrons originated from the same atom.Dative Covalent (or co-ordinate) bonding is where the shared electron pair originates from the same atom. Dative Covalent bonds behave like normal covalent bonds once formed

Requirements for Dative Covalent Bonding

1.An atom or an atom in a molecule or ion which has a lone pair of electrons available for donation. (often –ve charged)CALLED A NUCLEOPHILE

2.An atom or an atom in a molecule or ion which has an empty orbital which can accept an electron pair. (often +ve charged)

CALLED AN ELECTROPHILE

H+ and OH- example

•H+ ion (electronic configuration = 1s0)Has an empty s orbital to accept a lone pair

•OH- ion Has a lone pair which it can donate into the empty orbital

During the reaction, this happens to form a O-H dative covalent bond

O H H

Dative covalent bond

Boron compounds• All covalent compounds

• All have 6 electrons in outer shell of B

• All have an empty orbital

• All can accept an electron pairNN B

B

Shapes of Molecules•Valence Shell Electron Pair Repulsion theory (VSEPR theory)

•The shape of a molecule is dependent on the angle between the bonds in the molecule

•Electron Pairs will repel each other to the maximum extent

•Lone pairs repel to a greater extent than bonding pairs

•The bond angle (and hence shape) is therefore dependent on the number of bonding and lone pairs surrounds the central atom in the molecule

Working out shapes• Draw dot cross diagrams (or displayed formulae)• Work out the number of bonding and lone pairs around the central atom• Treat multiple bonds as 1 bonding pair

• H2

• BeI2 (covalent)• AlCl3 (covalent)• CF4 • H2S• PH3

• SiH4

• SO3

Molecule Number of electron Pairs

Max Distance repelled

Number of lone pairs

ActualBond angle

Shape

H2 1 180 0 180 Linear

BeI2 2 180 0 180 Linear

AlCl3 3 120 0 120 Trigonal Planar

CF4 4 109.5 0 109.5 Tetrahedral

H2S 4 109.5 2 104.5 Bent

PH3 4 109.5 1 107 Pyramidal

SiH4 4 109.5 0 109.5 Tetrahedral

SO3 3 120 0 120 Trigonal Planar

SF6 6 90 0 90 Octahedral

BeI2

CO2

AlCl3BH3

CF4

CH4

SO3

SF6

NF3

PH3

H2S

H2O

Complex molecules

• C2Cl4• CH3CH2OH

• Work out the bond angle (and hence shape around each central atom)

Intermolecular Forces

•What makes a substance a solid at RT?•Why is F2 a gas, while I2 is a solid?•How can you get liquid F2?•Why does H2O boil at 100oC while H2S boils at -60.3oC?

Intermolecular forces: THE FORCE OF ATTRACTION BETWEEN MOLECULES

Still a force of attraction between 2 oppositely charged entities

VAN DER WAALS FORCES

• Consider two F2 molecules

• What force of attraction could exist between them to enable F2 to liquefy?

VAN DER WAALS FORCES

• Electrons in a bond are not static

• Closer to one atom Temporary Dipole

• Causes dipole to form in neighbouring molecules Induced Dipoles

• Attraction between these two dipoles = VDW

Permanent Dipole- Dipole

• Molecules containing atoms of different electronegativity eg HCl, CH3Cl,

• Permanent dipoles are present

• Interactions occur between the δ+ve atom in one molecule and the δ-ve atom in the other molecule

Molecules without permanent dipoles when expected

• Symmetrical molecules

O=C=O

Dipoles cancel out

Hydrogen Bonding

• Especially large dipoles in molecules

• Especially large dipole-dipole attraction

• Called H-Bond

• Formed between molecules containing OH, NH or HF bonds

Drawing Hydrogen bonds

• Need at least 2 molecules

• Need to draw on dipoles on O and H atoms

• Need to draw lone pairs on O

• Draw straight line (often dotted) between lone pair on O to H on DIFFERENT Molecule

Water and DNA

WATER

• High MP than expected

• High surface tension

• Low density of ice

DNA

• Helical structure

• Strands held together

H-Bonds need to be broken

H-Bonds on surface pull molecules on surface in

H-Bonds hold Water molecules apart in Solid Water

Boiling Points•Non-Polar compoundsLOW as VDW weakBP increases as no of electrons in molecules increasesTemporary dipoles / induced dipoles increaseForce of attraction between molecules increase

•Polar compoundsHIGHER AS permanent dipole dipole attractions stronger than VDWNeed to look for a bond in the molecule containing two atoms of different electronegativities.

•Compounds containing OH, NH or HF bondsMuch higher as Hydrogen bonds much stronger than permanent dipole dipole

BUT• Intermolecular forces only found in simple molecular

structures and are MUCH WEAKER than the covalent, ionic or metallic bonds found in GIANT STRUCTURES

BP 0CNaCl 1250MgO 2390Cu 2600Carbon 4500

CH4 -79

CH3Cl -25

H2O 100

Giant Structures: BREAK STRONG CHEMICAL BONDS TO MELT

Simple Molecular Structures: BREAK Weak intermolecular forces to melt

In solid Methane Particle = CH4

In Gaseous Methane Particle = CH4

No Chemical bond is broken on melting / boiling