Electrons bonding and structure

Electrons bonding and structure

Module 2 aims

• Electron structure• Chemical bonding• Molecular shapes • Intermolecular forces• Bonding physical properties• Redox

Electron structure

ORBITALS

An orbital is... a region in space where one is likely to find an electron.Orbitals can hold up to two electrons as long as they have opposite spin; this is known as PAULI’S EXCLUSION PRINCIPAL.

Orbitals have different shapes...

ORBITAL SHAPE OCCURRENCE

s spherical one in every principal level p dumb-bell three in levels from 2 upwards d various five in levels from 3 upwards f various seven in levels from 4 upwards

An orbital is a 3-dimensional statistical shape showing where one is most likely to find an electron. Because, according to Heisenberg, you cannot say exactly where an electron is you are only able to say where it might be found.

DO NOT CONFUSE AN ORBITAL WITH AN ORBIT

Orbitals are not filled in numerical order because the principal energy levels get closer together as you get further from the nucleus. This results in overlap of sub levels. The first example occurs when the 4s orbital is filled before the 3d orbitals.

INC

REA

SIN

G E

NER

GY

/ DIS

TAN

CE

FRO

M N

UC

LEU

S

1 1s

22s2p

4s

33s3p3d

44p4d4f

PRINCIPAL ENERGY LEVELS

SUB LEVELS

1 1s

22s2p

3d

33s3p4s

44p4d4f

PRINCIPAL ENERGY LEVELS

SUB LEVELS

ORDER OF FILLING ORBITALS

THE FILLING ORDER

1s 2s 2p 3s 3p 3d 4s 4p 4d 4f 5s 5p 5d 5f 6s 6p 6d 7s 7p

HOW TO REMEMBER ...

1 1s

22s

2p

4s

3

3s

3p

3d

44p

4d

4f

INC

REA

SIN

G E

NER

GY

/ DIS

TAN

CE

FRO

M N

UC

LEU

S

This states that…“ELECTRONS ENTER THE

LOWEST AVAILABLE ENERGY LEVEL”

THE ‘AUFBAU’ PRINCIPAL

The following sequence will show the ‘building up’ of the electronic structures of the first 36 elements in the periodic table.Electrons are shown as half headed arrows and can spin in one of two directions

or

s orbitals p orbitals d orbitals

1 1s

22s

2p

4s

3

3s

3p

3d

44p

4d

4f

INC

REA

SIN

G E

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GY

/ DIS

TAN

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FRO

M N

UC

LEU

S

HELIUM

1s2

THE ELECTRONIC CONFIGURATIONS

Every orbital can contain 2 electrons, provided the electrons are spinning in opposite directions. This is based on...

PAULI’S EXCLUSION PRINCIPLE

The two electrons in a helium atom can both go in the 1s orbital.

‘Aufbau’Principle

1 1s

22s

2p

4s

3

3s

3p

3d

44p

4d

4f

INC

REA

SIN

G E

NER

GY

/ DIS

TAN

CE

FRO

M N

UC

LEU

STHE ELECTRONIC CONFIGURATIONS

OXYGEN

With all three orbitals half-filled, the eighth electron in an oxygen atom must now pair up with one of the electrons already there.

1s2 2s2 2p4

‘Aufbau’Principle

1 1s

22s

2p

4s

3

3s

3p

3d

44p

4d

4f

INC

REA

SIN

G E

NER

GY

/ DIS

TAN

CE

FRO

M N

UC

LEU


FLUORINE

The electrons continue to pair up with those in the half-filled orbitals.

1s2 2s2 2p5

1 1s

22s

2p

4s

3

3s

3p

3d

44p

4d

4f

INC

REA

SIN

G E

NER

GY

/ DIS

TAN

CE

FRO

M N

UC

LEU


NEON

The electrons continue to pair up with those in the half-filled orbitals. The 2p orbitals are now completely filled and so is the second principal energy level.In the older system of describing electronic configurations, this would have been written as 2,8.

1s2 2s2 2p6

1s1

1s2

1s2 2s1

1s2 2s2

1s2 2s2 2p1

1s2 2s2 2p2

1s2 2s2 2p3

1s2 2s2 2p4

1s2 2s2 2p5

1s2 2s2 2p6

1s2 2s2 2p6 3s1

1s2 2s2 2p6 3s2

1s2 2s2 2p6 3s2 3p1

1s2 2s2 2p6 3s2 3p2 1s2 2s2 2p6 3s2 3p3 1s2 2s2 2p6 3s2 3p4

1s2 2s2 2p6 3s2 3p5

1s2 2s2 2p6 3s2 3p6

1s2 2s2 2p6 3s2 3p6 4s1

1s2 2s2 2p6 3s2 3p6 4s2 1s2 2s2 2p6 3s2 3p6 4s2 3d1

1s2 2s2 2p6 3s2 3p6 4s2 3d2

1s2 2s2 2p6 3s2 3p6 4s2 3d3

1s2 2s2 2p6 3s2 3p6 4s1 3d5

1s2 2s2 2p6 3s2 3p6 4s2 3d5

1s2 2s2 2p6 3s2 3p6 4s2 3d6

1s2 2s2 2p6 3s2 3p6 4s2 3d7

1s2 2s2 2p6 3s2 3p6 4s2 3d8

1s2 2s2 2p6 3s2 3p6 4s1 3d10

1s2 2s2 2p6 3s2 3p6 4s2 3d10

HHeLiBeBCNOFNeNaMgAlSiPSClArKCaScTiVCrMnFeCoNiCuZn

ELECTRONIC CONFIGURATIONS OF ELEMENTS 1-30

ELECTRONIC CONFIGURATION OF IONS• Positive ions (cations) are formed by removing electrons from atoms• Negative ions (anions) are formed by adding electrons to atoms• Electrons are removed first from the highest occupied orbitals (EXC. transition metals)

SODIUM Na 1s2 2s2 2p6 3s1 1 electron removed from the 3s orbitalNa+ 1s2 2s2 2p6

CHLORINE Cl 1s2 2s2 2p6 3s2 3p5 1 electron added to the 3p orbitalCl¯ 1s2 2s2 2p6 3s2 3p6

FIRST ROW TRANSITION METALSDespite being of lower energy and being filled first, electrons in the 4s orbital are removed before any electrons in the 3d orbitals.

TITANIUM Ti 1s2 2s2 2p6 3s2 3p6 4s2 3d2

Ti+ 1s2 2s2 2p6 3s2 3p6 4s1 3d2

Ti2+ 1s2 2s2 2p6 3s2 3p6 3d2

Ti3+ 1s2 2s2 2p6 3s2 3p6 3d1

Ti4+ 1s2 2s2 2p6 3s2 3p6

Electronic structure summary

• Electrons occupy energy levels around the nucleus of the atom, where each shell has a principal quantum number.

• For principal quantum number, n = 1, the number of electrons is 2; for n = 2, the number is 8; then 18; then 32 electrons for n = 4.

• Main energy levels are sub-divided into sub-shells and these consist of orbitals called s, p and d-orbitals.

• Elements have an electronic configuration that can be shown in s, p or d notation, for example, sodium is 1s2, 2s2, 2p6, 3s1.

BONDING

The physical properties of a substance depend on its structure and type of bonding present. Bonding determines the type of structure.

TYPES OF BONDCHEMICAL ionic (or electrovalent)strong bonds covalent

dative covalent (or co-ordinate) metallic

PHYSICAL van der Waals‘ forces - weakestweak bonds dipole-dipole interaction

hydrogen bonds - strongest

THE IONIC BOND

Ionic bonds tend to be formed between elements whose atoms need to “lose” electrons to gain the nearest noble gas electronic configuration (n.g.e.c.) and those which need to gain electrons. The electrons are transferred from one atom to the other.

Sodium Chloride

Na ——> Na+ + e¯ and Cl + e¯ ——> Cl¯ 1s2 2s2 2p6 3s1 1s2 2s2 2p6 1s2 2s2 2p6 3s2 3p5 1s2 2s2 2p6 3s2 3p6

or 2,8,1 2,8 2,8,7 2,8,8

An electron is transferred from the 3s orbital of sodium to the 3p orbital of chlorine; both species end up with the electronic configuration of the nearest noble gas the resulting ions are held together in a crystal lattice by electrostatic attraction.

Physical properties of ionic compounds

Melting pointvery high A large amount of energy must be put in to overcome the

strong electrostatic attractions and separate the ions.

StrengthVery brittle Any dislocation leads to the layers moving and similar

ions being adjacent. The repulsion splits the crystal.

Electrical don’t conduct when solid - ions held strongly in the latticeconduct when molten or in aqueous solution - the ionsbecome mobile and conduction takes place.

Solubility Insoluble in non-polar solvents but soluble in waterWater is a polar solvent and stabilises the separated ions.

Much energy is needed to overcome the electrostatic attraction and separate the ions stability attained by being surrounded by polar water molecules compensates for this

IONIC BONDING

BRITTLE IONIC LATTICES

+ +

+ ++ +

+ +- -- -

- -

- -

+ +

+ +

IF YOU MOVE A LAYER OF IONS, YOU GET IONS OF THE SAME CHARGE NEXT TO EACH OTHER. THE LAYERS REPEL EACH OTHER AND THE CRYSTAL BREAKS UP.

IONIC COMPOUNDS - ELECTRICAL PROPERTIES

SOLID IONIC COMPOUNDS DO NOT CONDUCT ELECTRICITY

Na+Cl- Na+Cl-

Na+Cl-Na+ Cl-

Na+Cl- Na+Cl-

IONS ARE HELD STRONGLY TOGETHER

+ IONS CAN’T MOVE TO THE CATHODE

- IONS CAN’T MOVE TO THE ANODE

MOLTEN IONIC COMPOUNDS DO

CONDUCT ELECTRICITY

Na+ Cl-

Na+

Cl-

Na+

Cl-

Na+

Cl-

IONS HAVE MORE FREEDOM IN A LIQUID SO CAN MOVE TO THE ELECTRODES

SOLUTIONS OF IONIC COMPOUNDS IN

WATER DO CONDUCT ELECTRICITY

DISSOLVING AN IONIC COMPOUND IN WATER BREAKS UP THE STRUCTURE SO IONS ARE FREE TO MOVE TO THE ELECTRODES

Definition consists of a shared pair of electrons with one electron beingsupplied by each atom either side of the bond.compare this with dative covalent bonding

atoms are held togetherbecause their nuclei whichhave an overall positive chargeare attracted to the shared electrons

Formation between atoms of the same element N2, O2, diamond,graphite

between atoms of different elements CO2, SO2

on the RHS of the table;

when one of the elements is in the CCl4, SiCl4

middle of the table;

with head-of-the-group elements BeCl2

with high ionisation energies;

COVALENT BONDING

+ +

• atoms share electrons to get the nearest noble gas electronic configuration

• some don’t achieve an “octet” as they haven’t got enough electrons eg Al in AlCl3

• others share only some - if they share all they will exceed their “octet” eg NH3 and H2O

• atoms of elements in the 3rd period onwards can exceed their “octet” if they wish as they are not restricted to eight electrons in their “outer shell” eg PCl5 and SF6

COVALENT BONDING

Orbital theoryCovalent bonds are formed when orbitals, each containing one electron, overlap. This forms a region in space where an electron pair can be found; new molecular orbitals are formed.

SIMPLE MOLECULES

The greater the overlap the stronger the bond.

orbital containing 1

electron

orbital containing 1

electron

overlap of orbitals provides a region in space which can contain a pair of electrons

HYDROGEN

H H

Another hydrogen atom also needs one electron to

complete its outer shell

Hydrogen atom needs one electron to


atoms share a pair of electrons to form a single covalent bond

A hydrogen MOLECULE is formed

H H

HH

WAYS TO REPRESENT THE MOLECULE

PRESSING THE SPACE BAR WILL ACTIVATE EACH STEP OF THE ANIMATION

HYDROGEN CHLORIDE

Cl H

Hydrogen atom also needs one electron

to complete its outer shell

Chlorine atom needs one electron

to complete its outer shell

atoms share a pair of electrons to form a

single covalent bond

H Cl H Cl

WAYS TO REPRESENT THE MOLECULE


METHANE

CEach hydrogen atom needs 1

electron to complete its outer shell

A carbon atom needs 4 electrons to complete

its outer shell

Carbon shares all 4 of its electrons to form 4 single covalent bonds

H

H

H

H

H C H H

H

H C HH

H

WAYS TO REPRESENTTHE MOLECULE


WATER

OEach hydrogen

atom needs one electron to


Oxygen atom needs 2 electrons to complete its outer shell

Oxygen can only share 2 of its 6 electrons otherwise it will exceed the maximum of 82 LONE PAIRS REMAIN

H

H

H O H

H O

H

WAYS TO REPRESENTTHE MOLECULE


AMMONIA

NH H

H

NH

H H

H N HH

H N H

H

each atom needs one electron to complete

its outer shell

atom needs three electrons to complete

its outer shell

Nitrogen can only share 3 of its 5 electrons otherwise it will

exceed the maximum of 8A LONE PAIR REMAINS

WATER

OH

H OH

H

each atom needs one electron to complete

its outer shell

atom needs two electrons to complete

its outer shell

Oxygen can only share 2 of its 6 electrons otherwise it will exceed the maximum of 8

TWO LONE PAIRS REMAIN

H O H

H O

H

OXYGEN

O

each atom needs two electrons to complete its outer shell

each oxygen shares 2 of its electrons to form a

DOUBLE COVALENT BOND

O O O

O O

Bonding Atoms are joined together within the molecule by covalent bonds.

Electrical Don’t conduct electricity as they have no mobile ions or electrons

Solubility Tend to be more soluble in organic solvents than in water;some are hydrolysed

Boiling point Low - intermolecular forces (van der Waals’ forces) are weak; they increase as molecules get a larger surface area

e.g. CH4 -161°C C2H6 - 88°C C3H8 -42°C

as the intermolecular forces are weak, little energy is required toto separate molecules from each other so boiling points are low

some boiling points are higher than expected for a given massbecause you can get additional forces of attraction

SIMPLE COVALENT MOLECULES

A dative covalent bond differs from covalent bond only in its formationBoth electrons of the shared pair are provided by one species (donor) and it shares the electrons with the acceptor

Donor species will have lone pairs in their outer shells

Acceptor species will be short of their “octet” or maximum.

Lewis base a lone pair donorLewis acid a lone pair acceptor

DATIVE COVALENT (CO-ORDINATE) BONDING

Ammonium ion, NH4+

The lone pair on N is used to share with the hydrogen ion which needs two electrons to fill its outer shell.

The N now has a +ive charge as- it is now sharing rather than owning two electrons.

Boron trifluoride-ammonia NH3BF3

Boron has an incomplete shell in BF3 and can accept a share of a pair of electrons donated by ammonia. The B becomes -ive as it is now shares a pair of electrons (i.e. it is up one electron) it didn’t have before.

METALLIC BONDING

Involves a lattice of positive ions surrounded by delocalised electrons

Metal atoms achieve stability by “off-loading” electrons to attain the electronic structure of the nearest noble gas. These electrons join up to form a mobile cloud which prevents the newly-formed positive ions from flying apart due to repulsion between similar charges.

Atoms arrange in regular close packed 3-dimensional crystal lattices.

The outer shell electrons of each atom leave to join a mobile “cloud” or “sea” of electrons which can roam throughout the metal. The electron cloud binds the newly-formed positive ions together.

METALLIC BOND STRENGTHDepends on the number of outer electrons donated

to the cloud and the size of the metal atom/ion.

The strength of the metallic bonding in sodium is relatively weak because each atom donates one electron to the cloud.

The metallic bonding in potassium is weaker than in sodium because the resulting ion is larger and the electron cloud has a bigger volume to cover so is less effective at holding the ions together.

The metallic bonding in magnesium is stronger than in sodium because each atom has donated two electrons to the cloud. The greater the electron density holds the ions together more strongly.

Na

Mg

K

METALLIC PROPERTIES

MOBILE ELECTRON CLOUD ALLOWS THE CONDUCTION OF ELECTRICITY

For a substance to conduct electricity it must have mobile ions or electrons.

Because the ELECTRON CLOUD IS MOBILE, electrons are free to move throughout its structure. Electrons attracted to the positive end are replaced by those entering from the negative end.

Metals are excellent conductors of electricity

MALLEABLE CAN BE HAMMERED INTO SHEETS

DUCTILE CAN BE DRAWN INTO RODS AND WIRES

As the metal is beaten into another shape the delocalised electron cloud continues to bind the “ions” together.

Some metals, such as gold, can be hammered into sheets thin enough to be translucent.

METALLIC PROPERTIESMetals can have their shapes changed relatively easily

HIGH MELTING POINTS

Melting point is a measure of how easy it is to separate individual particles. In metals it is a measure of how strong the electron cloud holds the + ions.

The ease of separation of ions depends on the...

ELECTRON DENSITY OF THE CLOUDIONIC / ATOMIC SIZE

PERIODS Na (2,8,1) < Mg (2,8,2) < Al (2,8,3) m.pt 98°C 650°C 659°C b.pt 890°C 1110°C 2470°C

METALLIC PROPERTIES

Na+ Al3+Mg2+

MELTING POINT INCREASES ACROSS THE PERIODTHE ELECTRON CLOUD DENSITY INCREASES DUE TO THE GREATER NUMBER OF ELECTRONS DONATED PER ATOM. AS A RESULT THE IONS ARE HELD MORE STRONGLY.

HIGH MELTING POINTS

Melting point is a measure of how easy it is to separate individual particles. In metals it is a measure of how strong the electron cloud holds the + ions.

The ease of separation of ions depends on the...

ELECTRON DENSITY OF THE CLOUDIONIC / ATOMIC SIZE

GROUPS Li (2,1) < Na (2,8,1) < K (2,8,8,1) m.pt 181°C 98°C 63°C b.pt 1313°C 890°C 774°C

METALLIC PROPERTIES

MELTING POINT INCREASES DOWN A GROUPIONIC RADIUS INCREASES DOWN THE GROUP. AS THE IONS GET BIGGER THE ELECTRON CLOUD BECOMES LESS EFFECTIVE HOLDING THEM TOGETHER SO THEY ARE EASIER TO SEPARATE.

Na+ K+Li+

Chemical bonding summary

• Ionic bonding takes place when positive ions and negative ions are attracted in a giant ionic structure.

• Covalent bonding is the sharing of electron pair(s) between nuclei of atoms.

• The covalent bond and ionic bond are both very strong chemical bonds.

• A dative covalent bond is one formed in which both electrons are donated from the same atom.

REGULAR SHAPES

Molecules, or ions, possessing ONLY BOND PAIRS of electrons fit into a set of standard shapes. All the bond pair-bond pair repulsions are equal.All you need to do is to count up the number of bond pairs and chose one of the following examples...

C

2 LINEAR 180º BeCl2

3 TRIGONAL PLANAR 120º AlCl3

4 TETRAHEDRAL 109.5º CH4

5 TRIGONAL BIPYRAMIDAL 90º & 120º PCl5

6 OCTAHEDRAL 90º SF6

BOND BONDPAIRS SHAPE ANGLE(S) EXAMPLE

A covalent bond will repel another covalent bond

IRREGULAR SHAPES

If a molecule, or ion, has lone pairs on the central atom, the shapes are slightly distorted away from the regular shapes. This is because of the extra repulsion caused by the lone pairs.

BOND PAIR - BOND PAIR < LONE PAIR - BOND PAIR < LONE PAIR - LONE PAIR

OO O

As a result of the extra repulsion, bond angles tend to be slightly less as the bonds are squeezed together.

AMMONIA

ANGLE... 107°

SHAPE... PYRAMIDAL

HN NH H

HBOND PAIRS 3LONE PAIRS 1TOTAL PAIRS 4

H

H

N

H

H

H

N

H

107°H

H

N

H

• The shape is based on a tetrahedron but not all the repulsions are the same• LP-BP REPULSIONS > BP-BP REPULSIONS• The N-H bonds are pushed closer together• Lone pairs are not included in the shape

SHAPES OF IONS

N BOND PAIRS 3 PYRAMIDALLONE PAIRS 1 H-N-H 107°

BOND PAIRS 4 TETRAHEDRALLONE PAIRS 0 H-N-H 109.5°

N

H

H

H

N+H

H

H

H N+

BOND PAIRS 2 ANGULARLONE PAIRS 2 H-N-H 104.5°

N

H

H N

NH4+

NH2-

NH3

REVIEW

MOLECULES WITH DOUBLE BONDS

C O C OO

Carbon - needs four electrons to complete its shellOxygen - needs two electron to complete its shell

The atoms share two electronseach to form two double bonds

DOUBLE BOND PAIRS 2LONE PAIRS 0

BOND ANGLE...

SHAPE...

180°

LINEAR

O OC180°

Double bonds behave exactly as single bonds for repulsion purposes so the shape will be the same as a molecule with two single bonds and no lone pairs.

The shape of a compound with a double bond is calculated in the same way. A double bond repels other bonds as if it was single e.g. carbon dioxide

OTHER EXAMPLES

BrF5BOND PAIRS 5LONE PAIRS 1‘UMBRELLA’ANGLES 90° <90°

F

F F

FBr

F

FF FF Br

F

BrF3 BOND PAIRS 3LONE PAIRS 2’T’ SHAPEDANGLE <90°

F

F

Br

F

F

F

Br

F

SO42-

O

S

O-

O-

O

O S

O-

O-

OBOND PAIRS 4LONE PAIRS 0TETRAHEDRALANGLE 109.5°

ANSWERS ON NEXT PAGE

TEST QUESTIONS

For each of the following ions/molecules, state the number of bond pairsstate the number of lone pairsstate the bond angle(s)state, or draw, the shape

SiCl4

PCl6-

H2S

SiCl62-

PCl4+

BF3

TEST QUESTIONS

3 bp 0 lp 120º trigonal planar boron pairs up all 3 electrons inits outer shell

4 bp 0 lp 109.5º tetrahedral silicon pairs up all 4 electrons in

its outer shell 4 bp 0 lp 109.5º tetrahedral as ion is +, remove an electron

in the outer shell then pair up 6 bp 0 lp 90º octahedral as the ion is - , add one electron to

the 5 in the outer shell then pair up 6 bp 0 lp 90º octahedral as the ion is 2-, add two electrons

to the outer shell then pair up 2 bp 2 lp 92º angular sulphur pairs up 2 of its 6

electrons in its outer shell - 2 lone pairs are left

BF3

SiCl4

PCl6-

H2S

SiCl62-

PCl4+

ANSWER

For each of the following ions/molecules, state the number of bond pairsstate the number of lone pairsstate the bond angle(s)state, or draw, the shape

Shapes Number of bonded pairs of electrons

Number of lone pairs of electrons

Shape Bond angle

2 0 Linear 180

3 0 120

4 0 Tetrahedral

5 0 90 & 120

6 0 Octahedral

3 1 107

2 2 Angular

Shapes Number of bonded pairs of electrons

Number of lone pairs of electrons

Shape Bond angle

3 0 Linear 180

3 0 Trigonal planar 120

4 0 Tetrahedral 109.5

5 0 Trigonal-bipyramidal

90 & 120

6 0 Octahedral 90

3 1 Pyramidal 107

2 2 Angular 104

Molecular shapes summary• The shape of a molecule is determined by the repulsion

between bonded electrons and non-bonded electrons (lone pairs).

• Lone electron pairs repel more than bonded pairs of electrons and give rise to distorted shapes.

• By deducing the number of bonded electron pairs and lone pairs of electrons, the shape of a molecule may be predicted.

• BF3 is trigonal planar; CH4 and NH4+ are tetrahedral; SF6 is

octahedral; H2O is non-linear (V-shaped/bent); CO2 is linear and ammonia, NH3, as pyramidal

‘The ability of an atom to attract the electron pair in a covalent bond to itself’

Pauling Scale a scale for measuring electronegativityvalues increase across periodsvalues decrease down groupsfluorine has the highest value

H 2.1

Li Be B C N O F1.0 1.5 2.0 2.5 3.0 3.5 4.0

Na Mg Al Si P S Cl 0.9 1.2 1.5 1.8 2.1 2.5 3.0 K Br 0.8 2.8

ELECTRONEGATIVITY

INCREASEINC

REA

SE

Electronegativity

• No electronegativity difference between two atoms leads to a pure non-polar covalent bond.

• A small electronegativity difference leads to a polar covalent bond.

• A large electronegativity difference leads to an ionic bond.

Spectrum of bonding

100% covalent 100% ionic

0.0 3.2Difference in electronegativity

As a rough guide:• A difference greater than 2.0 will lead to essentially ionic bonding.• A difference that is less than 1.0 will lead to bonding that is mainly

covalent

Occurrence not all molecules containing polar bonds are polar overallif bond dipoles ‘cancel each other’ the molecule isn’t polarif there is a ‘net dipole’ the molecule will be polar

HYDROGEN CHLORIDE TETRACHLOROMETHANE WATER

POLAR MOLECULES

NET DIPOLE - POLAR NON-POLAR NET DIPOLE - POLAR

Factors affecting polarising ability

• Charge density of the positive ion• Increases as the positive ion gets smaller and

the charge density gets larger.• As a negative ion gets bigger it becomes easier

to polarise.

Type of dipole

Permanent

Dipole – dipole forces

Instantaneous Induced

Van der Waal forces

Although the bonding within molecules is strong, between molecules it is weak. Molecules and monatomic gases are subject to weak attractive forces.Instantaneous dipole-induced dipole forcesElectrons move quickly in orbitals, so their position isconstantly changing; at any given time they could beAnywhere in an atom. The possibility exists that one side has More electrons than the other. This will give rise to a dipole...

The dipole on one atom induces dipoles on others

Atoms are now attracted to each other by a weak forces

The greater the number of electrons, the stronger the attractionand the greater the energy needed to separate the particles.

NOBLE GASES ALKANES Electrons B pt. Electrons B pt.He 2 -269°C CH4 10 -161°CNe 10 -246°C C2H6 18 - 88°CAr 18 -186°C C3H8 26 - 42°CKr 36 -152°C

VAN DER WAALS’ FORCESINSTANTANEOUS DIPOLE-INDUCED DIPOLE FORCES

Occurrence occurs between molecules containing polar bondsacts in addition to the basic van der Waals’ forcesthe extra attraction between dipoles means thatmore energy must be put in to separate moleculesget higher boiling points than expected for a given mass

DIPOLE-DIPOLE INTERACTION

Mr °CCH4 16 -161SiH4 32 -117GeH4 77 -90SnH4 123 -50

NH3 17 -33PH3 34 -90AsH3 78 -55SbH3 125 -17

Mr °CH2O 18 +100H2S 34 -61H2Se 81 -40H2Te 130 -2

HF 20 +20HCl 36.5 -85HBr 81 -69HI 128 -35

Boiling pointsof hydrides

• an extension of dipole-dipole interaction

• gives rise to even higher boiling points

• bonds between H and the three most electronegative elements, F, O and N are extremely polar

• because of the small sizes of H, F, N and O the partial charges are concentrated in a small volume thus leading to a high charge density

• makes the intermolecular attractions greater and leads to even higher boiling points

HYDROGEN BONDING

Hydrogen bonds are formed by electronegative elements with a non-bonding pair of electrons available to bond with a H attached to an electronegative element.

HYDROGEN BONDING - ICE

each water molecule is hydrogen-bonded to 4others in a tetrahedral formation

ice has a “diamond-like” structure

volume is larger than the liquid making it

when ice melts, the structure collapsesslightly and the molecules come closer; theythen move a little further apart as they getmore energy as they warm up

this is why…a) water has a maximum density at 4°Cb) ice floats.

hydrogen bondinglone pair

HYDROGEN BONDING - HF

Hydrogen fluoride has a much higher boiling point than one would expect for a molecule with a relative molecular mass of 20

Fluorine has the highest electronegativity of all and is a small atom so the bonding with hydrogen is extremely polar

FH

FH

HF

HF

d +

d ¯

d +

d ¯

d +

d ¯d +

d ¯

hydrogen bonding

Intermolecular forces

• An intermolecular force exists between molecules and may include hydrogen bonding, dipole-dipole or van der Waals’ forces.

• Electronegativity is the ability of an atom in a covalent bond to attract a bonded pair of electrons towards itself.

• Hydrogen bonding arises in molecules in which a hydrogen atom is bonded to either an N or O atom.

• Water molecules, and other substances consisting of hydrogen bonding, have anomalous properties as a result.

Bonding and physical properties

• Metals consist of a close-packed arrangement of positive ions, through which delocalised electrons move.

• Metals are very good electrical conductors as a result of having mobile electrons.

• Giant structures have high melting and boiling points due to strong chemical bonds acting throughout the structure.

• Giant ionic structures conduct electricity when molten, and when dissolved in water due to mobile ions, not electrons.

Oxidation states

Element Oxidation state Unreacted elements 0 Na, K +1 Mg, Ca +2 Al +3 F -1 H Usually +1 O Usually -2 Cl Usually -1

The sum of the oxidation states in a neutral compound must be zero, and in any ion the sum of the oxidation states will be equal to the charge on the ion.

Working out oxidation states

• Enter zero for any uncombined element (e.g. Na or Cl2)• Put in the oxidation state(s) of the elements you are sure

of. (see above table)• Multiply this figure by the number of atoms of the

element that are present• Take this number change the sign and assign it to the

element unknown oxidation state.• Divide the number by the number of atoms of that

element in the compound• This final figure is the oxidation state of the element.

OXIDATION STATES

Q. Name the following... PbO2 lead(IV) oxide

SnCl2 tin(II) chloride

SbCl3 antimony(III) chloride

TiCl4 titanium(IV) chloride

BrF5 bromine(V) fluoride

manganese(IV) oxide shows that Mn is in the +4 oxidation state in MnO2

sulphur(VI) oxide for SO3 S is in the +6 oxidation state

dichromate(VI) for Cr2O72- Cr is in the +6 oxidation state

phosphorus(V) chloride for PCl5 P is in the +5 oxidation state

phosphorus(III) chloride for PCl3 P is in the +3 oxidation state

THE ROLE OF OXIDATION STATE IN NAMING SPECIES

To avoid ambiguity, the oxidation state is often included in the name of a species

Redox • An oxidation number indicates the formal charge of a

chemically combined particle in a compound.• The oxidation number of metals usually equals the group

number (as a positive value) and minus (8 – group number) for non-metals.

• An element has been oxidised if the oxidation number increases, and reduced if the oxidation number decreases.

• When they react, metals are normally oxidised (they lose electrons), whereas non-metals gain electrons and are reduced.

Documents

Electrons bonding and structure