ATOMIC STRUCTURE electron configuration of atoms and ions of the

ATOMIC STRUCTURE

electron configuration of atoms and ions of thefirst 36 elements (using s,p,d notation)

periodic trends in atomic radius, ionisationenergy and electronegativity

comparison of atomic and ionic radii

Electron Electron arrangements in the arrangements in the

atomatom

Boundary Surfaces of the Boundary Surfaces of the 4f4f Orbitals Orbitals

Energy levels, sub-levels and orbitalsEnergy levels, sub-levels and orbitals

Energy levels• Electrons occupy distinct energy levels (or

shells) arranged around the nucleus.

• The different main energy levels have different sub-levels in them

There are four types: s, p, d, f

Energy sub levels

• The number of sublevels is determined by the energy shell. One in Level 1, 2 in Level 2 etc

Orbitals

• Electrons occupy orbitals in pairs, with each member of the pair having opposite spin, which cancel.

• s sublevel has 1 orbital• p sublevel has 3 orbitals• d sublevel has 5 orbitals • f sublevel has 7 orbitals

Number of orbitals

A Cross Section of the Electron A Cross Section of the Electron Probability Distribution for a Probability Distribution for a 3p3p Orbital Orbital

1s

2s

2p 2p 2p

3s

3p 3p 3p

3d 3d 3d

Level 1 has just 1 sublevel consisting of one orbital (1s)

Level 2 has 2 sublevels (s and p) with one orbital in the first (2s) and 3 orbitals in the second (2p)

Level 3 has 3 sublevels (s,p and d) with one orbital in the first (3s), 3 orbitals in the second (3p)And 5 orbitals in the third (3d)

3d 3d

Energy Orbital arrangements in Atoms

Energylevel

1s

2s

2p

3s

3p

3d 4sn=4

n=3

n=2

n=1

Write the order of subshell filling using the graph1s 2s 2p 3s 3p 4s 3d

What atom is this?

Ne

Electrons fill orbit of lowest energy first.

Electron Arrangements in Orbitals

•In the normal ground state of an atom, the electrons occupy orbitals with the lowest possible energies. On heating, the electrons can be excited to orbitals with higher energy - the ‘excited state’.

•As the electrons fall back to lower energy levels (orbitals) they will emit electromagnetic radiation, which is often in the region of visible light ie. it appears coloured.

• Each element has its own characteristic emission spectrum that can be used to identify that element, its “chemical fingerprint”.

Atomic orbitals define regions of space in which there is a high probability of finding an electron.

Each orbital has a particular shape and associated energy values.

1.Each orbit can hold 2 electrons only an orbital containing two electrons is a filled orbit

2.Electrons filling the same orbit must have opposite spins indicated by arrows

3.Electrons fill the lowest energy sublevels first

4.The lowest or most stable arrangement of electrons in a sublevel is the one with the greatest number of parallel spins (Hunds Rule)

5.This means when orbitals of the same energy are available, electrons will avoid pairing if possible, by entering separate orbits

Rules For Filling Orbits

Energylevel

1s

2s

2p

3s

3p

3d 4sn=4

n=3

n=2

n=1

the electron configuration for Ne using the correct format is :

The Electron Arrangement for Ne is

1s2 2s2 2p6

Energylevel

1s

2s

2p

3s

3p

3d 4sn=4

n=3

n=2

n=1

Write the electron configuration then name the atom

Electronic Configuration - ways Electronic Configuration - ways to write themto write them

O: or

Cl: or

K: or

The electrons in the outermost shell of an atom are called its valence electrons

O: [He] 2s2 2p4

Core electrons Valence electrons

1s2 2s2 2p4 [He] 2s2 2p4

1s2 2s2 2p6 3s2 3p5 [Ne] 3s2 3p5

1s2 2s2 2p6 3s2 3p6 4s1 [Ar]4s1

Energylevel

1s

2s

2p

3s

3p

3d 4sn=4

n=3

n=2

n=1

[Ar] 4s2 3d2

Write the Electron Arrangement for Ti

Energy levels and sub-levelsEnergy levels and sub-levels

Main Main energy energy levellevel

Sub-Sub-levelslevels

Max. no. Max. no. of of

electron electron pairs in pairs in

sub-levelsub-level


electrons electrons in sub-in sub-levellevel


electrons electrons in main in main

levellevel






sub-levelsub-level





levellevel

11






sub-levelsub-level





levellevel11 ss






sub-levelsub-level





levellevel

11 ss 11






sub-levelsub-level





levellevel

11 ss 11 22






sub-levelsub-level





levellevel

11 ss 11 22 22






sub-levelsub-level





levellevel

11 ss 11 22 22

22






sub-levelsub-level





levellevel

11 ss 11 22 22

22ss 11 22






sub-levelsub-level





levellevel

11 ss 11 22 22

22ss 11 22

pp






sub-levelsub-level





levellevel

11 ss 11 22 22

22ss 11 22

pp 33






sub-levelsub-level





levellevel

11 ss 11 22 22

22ss 11 22

pp 33 66






sub-levelsub-level





levellevel

11 ss 11 22 22

22ss 11 22

88pp 33 66






sub-levelsub-level





levellevel

11 ss 11 22 22

22ss 11 22

88pp 33 66

33






sub-levelsub-level





levellevel

11 ss 11 22 22

22ss 11 22

88pp 33 66

33ss 11 22

pp 33 66






sub-levelsub-level





levellevel

11 ss 11 22 22

22ss 11 22

88pp 33 66

33

ss 11 22

pp 33 66

dd






sub-levelsub-level





levellevel

11 ss 11 22 22

22ss 11 22

88pp 33 66

33

ss 11 22

pp 33 66

dd 55






sub-levelsub-level





levellevel

11 ss 11 22 22

22ss 11 22

88pp 33 66

33

ss 11 22

pp 33 66

dd 55 1010






sub-levelsub-level





levellevel

11 ss 11 22 22

22ss 11 22

88pp 33 66

33

ss 11 22

1818pp 33 66

dd 55 1010






sub-levelsub-level





levellevel

11 ss 11 22 22

22ss 11 22

88pp 33 66

33

ss 11 22

1818pp 33 66

dd 55 1010

44






sub-levelsub-level





levellevel

11 ss 11 22 22

22ss 11 22

88pp 33 66

33

ss 11 22

1818pp 33 66

dd 55 1010

44

ss 11 22

pp 33 66

dd 55 1010




Max. no. of Max. no. of electron electron pairs in pairs in

sub-levelsub-level

Max. no. of Max. no. of electrons electrons in sub-in sub-levellevel

Max. no. of Max. no. of electrons electrons in main in main

levellevel

11 ss 11 22 22

22ss 11 22

88pp 33 66

33

ss 11 22

1818pp 33 66

dd 55 1010

44

ss 11 22

pp 33 66

dd 55 1010

ff






sub-levelsub-level





levellevel

11 ss 11 22 22

22ss 11 22

88pp 33 66

33

ss 11 22

1818pp 33 66

dd 55 1010

44

ss 11 22

pp 33 66

dd 55 1010

ff 77





sub-levelsub-level



levellevel

11 ss 11 22 22

22ss 11 22

88pp 33 66

33

ss 11 22

1818pp 33 66

dd 55 1010

44

ss 11 22

pp 33 66

dd 55 1010

ff 77 1414





sub-levelsub-level



levellevel

11 ss 11 22 22

22ss 11 22

88pp 33 66

33

ss 11 22

1818pp 33 66

dd 55 1010

44

ss 11 22

3232pp 33 66

dd 55 1010

ff 77 1414

Atomic OrbitalsAtomic Orbitals

Main energy Main energy levels, levels, nn

Energy Energy sub-sub-

levels, llevels, lOrbitalsOrbitals

Number of Number of electronselectrons

11 11 11ss 22

22 22 22ss22pp

22

66

33 33 33ss

33pp

33dd

22

66

1010

44 44 44ss

44pp

44dd

44ff

22

66

1010

1414

The order in The order in which the which the orbitals fill in orbitals fill in atomsatoms

Hydrogen 1s, Hydrogen 1s, 2s, and 3s 2s, and 3s OrbitalsOrbitals

Three Three 2p2p Orbitals Orbitals

Boundary Surfaces of the Boundary Surfaces of the 3d3d OrbitalsOrbitals

Energylevel

1s

2s

2p

3s

3p

3d 4sn=4

n=3

n=2

n=1

[Ar] 3d10 4s1

Write the Electron Arrangement for Cu

Copper has a unique electron configuration

Extra stability is gained when the 3d orbitals are half filled or completely filled

Energylevel

1s

2s

2p

3s

3p

3d 4sn=4

n=3

n=2

n=1

[Ar] 3d5 4s1

Write the Electron Arrangement for Cr Extra stability is gained when the 3d orbitals are half filled or completely filled

Chromium has a unique electron configuration

Orbitals Being Filled vs. Position in Periodic TableOrbitals Being Filled vs. Position in Periodic Table

Atomic and Ionic Radii

It is impossible to measure the atomic radius of an atom but we can measure the distances between adjacent nuclei in substances

In metals we look at half the distance between adjacent nuclei

In non metals we look at covalent bonds - that is half the distance between nuclei of like atoms covalently bonded together

Half this distance between nuclei

Ionic radiiRadii decrease across period

Radii increase down group

2006 Bonding Exam2006 Bonding ExamWrite the electron configuration for:Write the electron configuration for:

Cr Cr

MnMn

MnMn2+2+

Evidence AchievementAchievement with Merit

Two correct..Cr 1s22s22p63s23p64s13d5 OR [Ar] 4s13d5 OR [Ar]3d54s1

Mn 1s22s22p63s23p64s23d5

OR [Ar] 4s23d5 OR [Ar]3d54s2

Mn2+ 1s22s22p63s23p63d5 OR [Ar] 3d5

Trends in atomic and ionic radii

From your diagram for non transition elements note that:

Atomic and ionic radii increase going down a group because there are more electrons added as we move down a group

We surprisingly find radii decrease across a period .This is because as we go across a period the electron shells are pulled closer into the nucleus

Why are they pulled closer. Bridget?

Because the number of electrons in the outer shell and the number of protons in the nucleus also increase.

All the electrons in the shell are about the same distance from the nucleus, so adding extra electrons isn’t going to make much difference

But as we increase the number of protons in the nucleus the electrostatic attraction from the nucleus becomes greater on the outer shell electrons pulling them closer

Why do radii decrease moving across the period (row)

Each atom in the group has the same number of valence electrons and the same effective nuclear charge from the nucleus on these valence electrons.

But there are more electron shells as we go down the group.

Because the outer shells are protected or shielded from the nucleus by the inner electrons the valence shells end up further away from the nucleus and therefore the electrostatic attraction between the valence electrons and the nucleus decreases

Why the atomic Radii increase down a group

Ionic Radii

Ionic radii of negative ions are larger than their corresponding atom

As more electrons are added to the same number of protons repulsion pushes the valence electrons further apart.

The more negative the ion the larger it is eg N 3-

Ionic Radii

Ionic radii of positive ions are smaller than the corresponding atom

Positively charged ions have lost electrons, but the number of protons remains the same ... so the remaining electrons are held more strongly than before

The more positive the ion the smaller it is eg Al 3+

2006 Bonding Exam2006 Bonding Exam Compare the relative sizes of the CaCompare the relative sizes of the Ca2+2+ and Cl and Cl––

ions, and explain the difference in their radii.ions, and explain the difference in their radii.

Evidence AchievementAchievement with Merit

Cl– is larger or Ca+2 smaller.

Both have same number of shells or electron arrangement. Ca2+ has more protons or nuclear charge is greater, so the electrostatic attraction between the valence electrons and the nucleus is stronger, making Ca2+ smaller.

Correct size. Size and explanation correct.

Emission spectraof a hydrogen electron

http://en.wikipedia.org/wiki/File:Emission_spectrum-H.png

Let us begin by examining the figure on the left that represents the 'spectrum' of the hydrogen atom. In this spectrum, each line represents a state or 'energy level' in which the hydrogen electron can exist, and the arrows indicate possible transitions between these energy levels.

When a hydrogen electron in one energy level makes a transition to a lower energy level, it emits energy in the form of photons, and the wavelength (or colour) of the emitted light is completely determined by the energy difference between the two levels. The spectrum of light shown below arises because the atom 'falls' from various higher states to the second lowest state. The lowest state is called the ground state.

Ionisation Energies

Define the term “Ionisation Energy” “The energy required to remove the

least tightly held electron from each atom of one mole of gaseous atoms or ions”

Do you think ionisation energies are always endo or exothermic Hannah?

Ionisation Energies

Ionisation energy is measured in kJ mol-1

The first ionisation energy of an atom is the energy to remove the 1st electron from the atom

The second ionisation energy removes the second electron

The first ionisation energy of Neon is +2087 kJmol-1

The thermochemical equation for this is written as:

1. Ne (g) Ne+ (g) + e H = +2087 kJmol-1

The second ionisation energy of Neon is +6128 kJmol-1

2. Ne+ (g) Ne2+ (g) + e H = +6128 kJmol-1

The second ionisation energy is higher because an electron is being removed from a positive ionWhy is the second ionisation energy always larger

than the first IE ? Daniel

We can tell the electronic structure of a K atom from its successive ionisation energies

Can you see how?

Remember the ionisation exercise last term – turn to it please

You have graphed the ionisation energies

now lets look at the questions

Recognising Periodic Ionisation Trends

As you go down a group the ionisation energies go down

eg He, Ne , Ar, Kr

The general trend across a period is of increasing ionisation energies eg Na to Ar

The zig zag shape of the graph indicates some atoms are more stable than others

The first ionisation energy of aluminium is 578 kJmol-1

Write the thermodynamic equation for this data:1. Al (g) Al+ (g) + e H = 578 kJmol-1

2. Explain why the first ionisation energy of :

a. Argon is lower than that of neon

b. Lithium is lower than that of neon

The valence electrons of argon are further away from the nucleus than those of neon, so it takes less energy to remove one from argon than neon

Lithium’s nuclear charge is +3, while neon’s is +10, so neon’s valence electrons are held more strongly than lithium’s

Answer these

Extension questions – Ionisation Energy Trends

Electron Configurations beyond Ar

The 4s orbital has slightly lower energy than the 3d orbitals, so the 4s orbital is filled before the 3d orbital

3d

K

Ca

4s

[Ar] 4s1 [Ar]

4s 3d

[Ar] [Ar] 4s2

Electron Configurations Of Transition metals

The 10 elements beyond Ca are called the transition metals and are found in the middle of the periodic table, almost all (except Cu and Cr) have 2 electrons in the 4s orbital

Electron Configurations Of Transition metals 3d

Sc

4s

[Ar] 3d1 4s2 [Ar]

Ti [Ar] [Ar] 3d2 4s2

V [Ar] [Ar] 3d3 4s2

Cr [Ar] [Ar] 3d5 4s1

Mn [Ar] [Ar] 3d5 4s2

Fe [Ar] [Ar] 3d6 4s2

Co [Ar] [Ar] 3d7 4s2

Ni [Ar] [Ar] 3d8 4s2

Cu [Ar] [Ar] 3d10 4s1

Zn [Ar] [Ar] 3d10 4s2

Write the electron configurations for the following

a. 7N3-

b. 33As

c. 28Ni2+

d. 24Cr

e. 35Br -

f. 12Mg2+

1s2 2s2 2p6

[Ar] 3d10 4s24p3

[Ar] 3d8

[Ar] 3d5 4s1

[Ar] 3d10 4s24p6

1s2 2s2 2p6

StarterReferring to your graph of atom radii (previous slide)

Why do atomic radii decrease as we move across a period?

As we move across a period an extra electron is added to the same shell while at the same time an extra proton is added to the nucleus.

The added electron is about the same distance away from the nucleus as the other outer shell electrons.

This means that there is a stronger electrostatic attraction from the nucleus on these outer shell electrons these electrons are then pulled closer to nucleus reducing the atom radius.

Referring to your graph of atom radii Why do atomic radii increase as we move

down a group? Give points

As we moving down a group an extra electron is added to a *different shell (further from the nucleus) while at the same time an extra proton is added to the nucleus.*most important

The added electron is shielded from the nucleus by the inner electrons.

Because of this the electrostatic force between the valence electrons and the nucleus decreases and the valence shells end up further away from the nucleus creating a larger radius

Remember the 4s orbital fills first and also empties first

Variable Oxidation states of MnVariable Oxidation states of Mn

Use the expt sheet Use the expt sheet

Transition Metal Colours

Transition MetalsTransition metals have varying oxidation states by holding on to or losing varying amounts of d electrons

Eg Fe2+ , Fe 3+

Cu +, Cu 2+

V 2+ , V 3+ , V 4+ , V 5+

Most transition metals have partially filled d orbitalsAnd can form coordination complexes with other species eg H2O, Cl , NH3 to name a few

When transition metal ions form diative bonds with the above species coloured complexes occur.

Transition Metal PropertiesTransition metals have variable _______ ______and most form ________ complexes eg The [Ti(H2O)6]3+ complex appears mauve because yellow green light is absorbed and a combination of visible light is reflected or transmitted to give a mauve colour

Another transition metal c_____ forms a complex ion with 6 H2O to give a ________ blue complex – name and give the formula of another Cu complexThese properties are due to the partially filled d subshell

opper[Cu(H2O)6] 2+

colouredoxidation states

[Cu(NH3)4] 2+ copper tetra amine

Transition Metal complex ions

They also form other complex ions in which the central metal ion is surrounded by other neutral molecules such as H2O or NH3, or by anions, such as Cl or OH. These are called ligands.

The non-bonded pairs of electrons on the ligands are donated to the metal cation and form co-ordinate (or dative) covalent bonds to the central metal ion.

This is all due to a partially filled ‘d’ shell

Transition Metal complex ions

The number of ligands attached to the central atom is called the co-ordination number.

For transition metals the most common co-ordination number is 6 and the complex ion has an octahedral arrangement.

[Fe(H2O)6]3+ [Co(NH3)6]

3+

cobalt(II) nitrate, Co(NO3)2 (red);

potassium dichromate, K2Cr2O7 (orange);

potassium chromate, K2CrO4

(yellow)

nickel(II) chloride NiCl2 (green)

copper(II) sulfate,

potassium permanganate, KMnO4 (purple).

CuSO4 5H2O

http://upload.wikimedia.org/wikipedia/commons/5/57/Coloured-transition-metal-solutions.jpg

cobalt(II) nitrate,

K2Cr2O7

K2CrO4

nickel(II) chloride NiCl2

CuSO4 5H2O

potassium permanganate, KMnO4

.The colour we see is the light that is reflected after it is absorbed to elevate the electronsEg a hydrated Cu ion absorbs orange light and reflects blue

The 3d sub shell is split and allows for electrons to be elevated to orbits within the subshell often absorbing light in the coloured region.

http://upload.wikimedia.org/wikipedia/commons/5/57/Coloured-transition-metal-solutions.jpg

The number of oxidation states generally increases with the number of unpaired electrons, with elements in the middle of the row (eg. Mn) having the widest range of oxidation states. The maximum oxidation state for transition elements up to Mn is equal to the total number of valence electrons available.

V Cr Mn Fe Cu 3d 34s 2 3d 54s 1 3d 54s 2 3d 64s 2 3d104s 1

+2 to +5 +2 to +6 +2 to +7 +2, +3 +1, +2

Oxidation states

Oxidation States of Transition metals

Transition metal ions with a high charge density such as Fe3+ form aqueous solutions that are acidic. This is due to the reaction of the aquo complex ion with water, as shown below.

[Fe(H2O)6]3+ + H2O [Fe(H2O)5OH]2+ + H3O

+

yellow

Note that many complex ions contain twice as many ligands as the charge on the central atom eg [Cu(NH3)4]

2+ or [Ag(NH3)2]+

Zn(OH)2 + 2OH [Zn(OH)4] 2

white ppt colourless soln

Cu(OH)2 + 4NH3 [Cu(NH3)4 ] 2 + + 2OH

light blue ppt dark blue soln

Other complex ions with ammonia are [Zn(NH3)4 ] 2+

and [Ag(NH3)2 ] +.

[Cu(H2O)6 ] 2+ + 4Cl [CuCl4 ]

2 + 6H2O light blue yellow

Oxides/hydroxides of MetalsOxides/hydroxides of Metals The alkali and alkaline metal oxides form basic solutionsThe alkali and alkaline metal oxides form basic solutions

eg Naeg Na22O + HO + H22O 2NaOHO 2NaOH

Transition metal oxides/hydroxides - as TM s increase their Transition metal oxides/hydroxides - as TM s increase their oxidation state oxidation state they become more acidic egthey become more acidic eg

Eg MnEg Mn22OO77 + H + H22O 2HMnOO 2HMnO44

What’s the ON of Mn in MnWhat’s the ON of Mn in Mn22OO77 ? ?

Some transition metal hydoxides are amphoteric (react as acids or bases)Some transition metal hydoxides are amphoteric (react as acids or bases)

Cr(OH)Cr(OH)33 + 3H + 3H++ Cr Cr3+ 3+ + 3H+ 3H22OO

Cr(OH)Cr(OH)33 + OH + OH-- [Cr(OH) [Cr(OH)44]]--

BondingBonding, ,

All noble gases have a stable electron configuration

Ar

Ne; 1s2, 2s2, 2p6

Ar; [Ne] 3s2, 3p6

Only Only ValenceValence Electrons are Electrons are involved in Chemical Bondinginvolved in Chemical Bonding

Lewis Dot Structure = Valence e-s are represented by dots about the elemental symbol

Lewis Symbols

Represent the number of valence electrons as dots

Valence number is the same as the Periodic Table Group No

H

Li Be B C N O F Ne

He

Na; Is2, 2s2, 2p6, 3s1 = [Ne] 3s1

Lewis Structure = Na

For example,

Groups 1 2 3 4 5 6 7 8

Summary of main structures of solidsSummary of main structures of solids

StructureStructureParticles Particles

in the in the SolidSolid

Bonds Bonds between between

the the ParticlesParticles

Boiling Boiling PointPoint

Solubility Solubility in Waterin Water

Electrical Electrical ConductivityConductivity ExamplesExamples

Molecular

small covalent

moleculeslow

Soluble/insoluble depends

on polarity

Giant covalent network

high insoluble

Ionicstrong ionic

bondshigh

sodium chloride,

magnesium oxide

Metallicstrong

metallic bonds

high insoluble

magnesium, iron,

copper, sodium

atoms

very weak forces

Doesn’t conduct In any state methane,

iodine,

water

Strongcovalent bonds

Doesn’t conductin any state

diamond, silicon dioxide

positive and

negative

ions

soluble

conducts whenMolten or in Solution but not when solid

positive ions in a sea of electrons

conducts when

solid and when

molten

PCl Cl

ClPCl Cl

Cl

Rules for Drawing Lewis Structures

• First add up the number of valence electrons from each atom, add one for each negative charge, subtract one for each positive charge.

• The central atom is usually written first in the formula• Complete the octets of atoms bonded to the central atom

(remember that H can only have two electrons)• Place any left over electrons on the central atom, even if

doing so it results in more than an octet• If there are not enough electrons to give the central atom

an octet , try multiple bonds

E.g. 1. PCl3Total Number of valence electrons = 5 + (3 x 7) = 26

PCl Cl

Cl

Draw the following lewis structures

O O x

xxx

xx

Oxygen gas O2

Nitrogen gas N2

H H x

Hydrogen gas H2

N N x x

xxx

Ammonia NH3

N H x

xx

H

H

Draw the following lewis structures

Phosphorus Trichloride PCl3

P Cl x

xx x

xx

x x

Cl x x

x x

x xx

Cl x x

x xx x

C Cl x

xx x

xx

x x

Cl x x

x x

x xx

Cl x x

x xx x

H

Chloromethane CHCl3

(chloroform)

Exceptions to the Octet ruleExceptions to the Octet rule

•An atom can have fewer than 8 e - eg B only needs 6 valence electrons and Be needs 4 valence electrons

•Second row elements never exceed the octet rule. (C, N, O, and F always obey the octet rule)

•Third row elements often satisfy octet rule, but can exceed it by using empty d orbitals

P Cl

Cl

Cl

Cl

Cl

:: :

:

::

::

::: :

: :

: ::

:: :

Elements beyond the second row of the periodic table (Na and beyond) are able to hold more than 8 electrons in their valence shell, due to the presence of d-orbitals.

The maximum number of electrons any atom can have is double its original number (each electron may invite one ‘friend’). Thus S can have a maximum of 12 electrons around it, but B can only have 6.

Draw the Lewis diagram for I3

–.Each I atom has 7 valence electrons. Include one more electron for the minus charge, and the total is 22 electrons.

II I::

Count electrons

Place atoms and charge.

Electrons where there must be bonds.

Make outer atoms happy.

Place rest of electrons on central atom.

All atoms satisfied.

: :::

::

:: : –

The middle I atom can accommodate 10 electrons because of an expandedd subshell

E.g. 2; CHBr3

Total Number of valence electrons = 4 + 1 + (3 x 7) = 26

BrCBr HBr

Exceptions to the Octet Rule in Covalent Bonding

1. Molecules with an odd number of electrons

2. Molecules in which an atom has less than an octet

3. Molecules in which an atom has more than an octet

1. Odd Number of Electrons

NO Number of valence electrons = 11

N O N O Resonace Arrows

NO2 Number of valence electrons = 17

Resonance occurs when more than one valid Lewis structure can be written for a particular molecule (i.e. rearrange electrons)

N OO N OO N OO

O O O OOxygen is a ground state"diradical"

O2

Molecules and atoms which are neutral (contain no formal charge) and with an unpaired electron are called Radicals

Exercise: Draw Lewis diagrams of the following molecules.

CH4, H2S,

CH H

H

H

S

H

H

C OH

HPCl

Cl

Cl C SSC O

H

H

H

H

H2CO, PCl3, CS2, CH3OH

We should be able to reach the same conclusion using the Periodic Table,

Cl is furthest to the right and to the top of the Periodic Table, so is the most electronegative. Se is furthest to the left (‘metallic like’) and towards the bottom. Therefore, difference in electronegativity should be the greatest!

Home work due Tuesday - read unit 4 page 16 complete Q 2,3, 4 and 5 on page 18

StarterStarter

Electronegativity is the ___________ Electronegativity is the ___________ an atom has for _________ in a bond an atom has for _________ in a bond

3. More than an Octet

PCl5

Elements from the third Period and beyond, have ns, np and unfilled nd orbitals which can be used in bonding

P : (Ne) 3s2 3p3 3d0

Number of valence electrons = 5 + (5 x 7) = 40

P

Cl

ClCl

ClCl

10 electrons around the phosphorus

2. Less than an Octet

Includes halides of B, Al and compounds of Be

Group 3A atom only has six electrons around it

B

Cl

Cl Cl

Starter - Draw the lewis diagram of the BCl3 moleculeGive the :

shape bond anglePolarity

Triangular planar

120 o

Non polar

3. More than an Octet

SF4

S : (Ne) 3s2 3p4 3d0

Number of valence electrons = 6 + (4 x 7) = 34

SF

F

F

F

The Larger the central atom, the more atoms you can bond to it – usually small atoms such as F, Cl and O allow central atoms such as P and S to expand their valency.

AlCl3

In the solid form Aluminium chloride is an ionic solid of AlCl3

However, it sublimes at 192 °C to a vapour of Al2Cl6 molecules

AlCl

Cl Cl

ClAl

Cl

Cl

3s2

Al

3p1

[Ne] Ground state

3s2

Al

3p1

[Ne] Promotion of e-

2 Cl electrons can form covalent bond

B2H6

A Lewis structure cannot be written for diborane. This is explained by a three-centre bond – single electron is delocalized over a B-H-B

BH

H H

HB

H

H

COVALENT BONDINGElemental hydrogen exists as a diatomic

molecule H + H H2

• Each hydrogen has one electron in the valence shell.

:H H H H

• However, both hydrogen atoms have the same number of electrons and the same +1 nucleus. The equal pull on the same number of electrons results in a sharing of electron density.

• What are we referring to here Ethiopia ?

ElectronegativityElectronegativity

Electronegativity (EN)Electronegativity (EN) = the ability of a bonding = the ability of a bonding atom to pull electrons from another atomatom to pull electrons from another atom

The higher the EN, the greater the pullThe higher the EN, the greater the pull Fluorine is most electronegative, EN decreases Fluorine is most electronegative, EN decreases

as you move awayas you move away Comparing the EN of two atoms tells you if the Comparing the EN of two atoms tells you if the

electrons are evenly shared (non polar electrons are evenly shared (non polar covalent) or unevenly (polar covalent). covalent) or unevenly (polar covalent).

Electronegativity is defined as the ability of an atom in a molecule to attract electrons to itself

Prof. Linus PaulingNobel Prize for Chemistry 1954Nobel Prize for Peace 1962

Electronegativity is a function of two properties of isolated atoms;The atom’s ionization energy (how strongly an atom holds onto its own electrons)The atom’s electron affinity (how strongly the atom attracts other electrons)

Figure 4.2

Remember that EN(F) is the highest.

Increasing EN

Polar BondsPolar Bonds

When the electronegativities of H and X differ, the resulting bond is polar:

H—X +

The bond has partial ionic character

The three major types of intramolecular bond can be described by the electronegativity difference:

Non-Polar Covalent – Bonds which occur between

atoms with little or no electronegativity difference (less

than 0.5).

Polar Covalent – Bonds which occur between atoms

with a definite electronegativity difference (between 0.5

and 2.0).

Ionic – Bonds which occur between atoms with a large

electronegativity difference (2.0 or greater), where

electron transfer can occur.

E.g. F-F (2.5 – 2.5 = 0) is non-polar covalent

H-F (4.0 – 2.1 = 1.9) is polar covalent

LiF (4.0 – 1.0 = 3.0) is ionic

H F+ -

Dipole moment

Examining the formation of NaClNa + Cl NaCl

IONIC BONDING

Sodium has alow ionization energyit readily loses this electron .When sodium loses the electron, it becomesisoelectronic with Ne.

Na Na+ + e-

Chlorine has a high ionization energy.

When chlorine gains an electron, it becomes isoelectronic with Ar.

:

..

..Cl: e

..

..Cl:

Two viewpoints of NaClTwo viewpoints of NaCl

You must also know and justify the differences betweenIonic solidsMolecular solidsCovalent network solidsmetals

In terms of: boiling point/melting point – ie justify using the forces between particlesConductivitySolubility

structure and bonding

The Effect of an The Effect of an Electric Field on Electric Field on Hydrogen Fluoride Hydrogen Fluoride MoleculesMolecules

Draw the Lewis diagram for NO3–.

Total number of electrons:

5 + 3(6) + 1 = 24.:

All 24 electrons have been distributed. The O atoms have 8 electrons each, which satisfies them, but the N atom only has 6 electrons around it, which is unacceptable.

NO O

O

: :

: :

::

:::

:

:–

All atoms now have an acceptable number of electrons. This is the final Lewis diagram.Move one pair of

electrons to make a double bond.

Bestchoice homework

Particle Properties 3.4 Lewis ex first Then VSEPR :

Problem set 1

Problem set 2

Three Possible Types of BondsThree Possible Types of Bonds

(a) Pure covalent

(b) Polar Covalent

(c) Ionic

Type of BondType of Bond Electron Electron DistributionDistribution

Examples Examples of Bondsof Bonds

Physical Physical PropertiesProperties

IonicIonic(Very different (Very different ENs)ENs)

Transfer to form Transfer to form ions (beyond ions (beyond polar)polar)

LiF, NaClLiF, NaCl Crystalline solids Crystalline solids with high mpswith high mps

Polar CovalentPolar Covalent(Differing ENs)(Differing ENs)

Uneven SharingUneven Sharing HCl, HHCl, H22O, O,

SiClSiCl44, NH, NH

33

gases, liquids or gases, liquids or solids with lower solids with lower mpsmps

Nonpolar Nonpolar CovalentCovalent(similar ENs)(similar ENs)

Even SharingEven Sharing molecular molecular elements, C-H elements, C-H bonds bonds (exception)(exception)

gases, liquids and gases, liquids and solids with very solids with very low mpslow mps

Dipole Moment occurs in any polar covalent bond, because of an unequal sharing of the electron pair between two atoms

E.g. Which of the following bonds is most polar: S-Cl,

S-Br, Se-Cl or Se-Br?S-Cl (3.0 – 2.5) = 0.5S-Br (2.8-2.5) = 0.3Se-Cl (3.0-2.4) = 0.6Se-Br (2.8-2.4) = 0.4

Therefore, Se-Cl is the most polar!

Shapes of moleculesShapes of molecules

Valence-Shell Electron-Pair Repulsion Theory (VSEPR)

In molecules there are 2 types of electron

1. Bonding Pairs2. Non-bonding or lone pairs The combinations of these determine the shape of the molecule The outer pairs of electrons around a covalently bonded atom minimize repulsions between them by moving as far apart as possible

Shapes of molecules

The shape of a molecule is determined by the positions of atoms within that molecule.

The position of those atoms is determined by the arrangement of the electron sets (both bonding and non bonding sets) around the central atom.

Electron sets want to repel each other as much as possible

If there are any non-bonding (‘lone pair’) electrons around the central atom, they will squeeze the bonding pairs (with their attached atoms) closer together.Lone pairs are closer to the central atom than bonding pairs (because they are not shared between two nuclei). They push the bonding pairs slightly closer together than if all pairs were bonding. This is known as lone pair repulsion.

The shape of the molecule depends on: 1. the number of atoms linked to the central atom

2. the total number of electron sets (bonded and non-bonded electron pairs) around the central atom.

NOTE: Multiple electron pairs (i.e. double and triple bonds) are considered to be only one electron set.

The shapes of molecules are determined by the way clouds of electrons are arranged around the central atom in the molecule.

A molecule containing only a single cloud of electrons must be linear.

H—H

Arrangements of electron clouds

Two clouds arrange themselves on opposite sides of the central atom.

The bond angle will be 180°.

O = C = O

Since clouds of electrons are negatively-charged, they repel each other.

Two clouds

Notice that the double bonds in CO2 each act as a single cloud of electrons.

Adding a third cloud of electrons will change the bond angle from 180° to 120°.

All the atoms still lie on a flat plane (like a sheet of paper). The shape is trigonal planar.

Three clouds

When a fourth cloud is added, the previous clouds are pushed downwards.

This shape is tetrahedral.

The bond angle is now 109°.

HH

H

H

C

Four clouds

Remember this number. A circle divided in 4 makes an angle of 90°, but a sphere divided in 4 makes an angle of 109°.

A fifth cloud enters from below, raising the lower 3 clouds back to a flat triangle.

This shape is trigonal bipyramid.

Bond angles around the central triangle are 120°, while the angle between the triangle and the top and bottom atoms is 90°.

P

ClClCl

Cl

Cl

Five clouds

A sixth cloud squeezes around the centre and making all bond angles 90°.

S

FF

FF

F

F

This shape is octahedral.

Six clouds

Three clouds of electrons around central atom:

Trigonal planar

Bent Linear

Bond angle = 120°

Bond angle = 119°

OSO

O SOO

Four clouds of electrons around central atom:

H

HH

HCH

HHN

HH

OH

Cl

Bond angle = 109°

Bond angle = 107°

Bond angle = 105°

Tetrahedral Trigonal pyramid

Bent Linear

The N atom lies above the plane of the H atom.Pyramid with an equilateral triangle as the base.

Trigonal Pyramidal

Five clouds of electronsWhen five clouds of electrons are around the central atom there are two different bond angles — 120° and 90°. Any lone pairs will be positioned around the central triangle first, where they can be 120° away from other clouds.

Trigonal bipyramid

See-saw T-shaped Linear

P F

F

F

FF

S F

F

FF

Cl F

F

F

I

I

I

Six clouds of electrons

Octahedral Square pyramid Square planar

SFFF

F

F

F

BrFFF

F

F

XeFFF

F

Polar Molecules = Molecules with permanent dipole moments

HCl has only one covalent bond (which is polar). Therefore, its dipole moment = H-Cl bond dipole

In a molecule with two or more polar bonds, each bond has a dipole moment contribution = bond dipole

Net dipole moment = vector sum of its bond dipoles

Linear Molecules: CO2 is Non-polar C OO

Because CO2 dipoles are orientated in opposite directions.The dipoles have equal magnitudes and therefore they cancel

Net dipole = 0

The shape of the molecule and

Number of e- sets

Are there polar bonds?

Are polar bonds arranged symmetrically?

Will polar bonds cancel each other?

Will there be a net dipole on the molecule?

ClCl

ClC

Cl The CCl4 molecule is a tetrahedral shape because the 4 bonding electron sets repel each other evenly

The CCl4 molecule has 4 polar C-Cl bonds

These 4 polar C-Cl bonds are arranged symmetrically in the CCl4 molecule

Because the 4 polar C-Cl bonds are arranged symmetrically in the CCl4 molecule they cancel each other and the CCl4 has no net dipole moment and is nonpolar

The shape of the molecule

And number of electron

Sets around central atom

Are there polar bonds?

Are polar bonds arranged symmetrically?

Will polar bonds cancel each other?

Will there be a net dipole on the molecule?

HH

HN

The NH3 molecule has 4 electron sets (3 bonding and 1 non bonding) around the central N atom which repel each other. Because the overall shape of the NH3 molecule is determined by the position of the atoms it is a trigonal pyramid shape

The NH3 molecule has 3 polar N-H bonds

These polar N-H bonds are arranged unsymmetrically in the NH3 molecule

Because the 3 polar N-H bonds are arranged unsymmetrically in the NH3 molecule they do not cancel each other and the NH3 has an overall net dipole moment causing the NH3 to be polar

x x

Intermolecular Forces:

These are generally much weaker than covalent or ionic bonds.

Less energy is thus required to vaporize a liquid or melt a solid.

Boiling points can be used to reflect the strengths of intermolecular forces (the higher the Bpt, the stronger the forces)

Weak Intermolecular ForcesWeak Intermolecular Forces

Consist of:

• Temporary dipole – temporary dipole interactions

•Dipole-dipole interactions

These exist in Polar Molecules

•The permanent dipole interacts with other permanent dipoles to cause attraction.

•exist between all polar molecules,

•The strength of the attraction depends on the polarity of the molecule.

Permanent dipolesPermanent dipoles

Electrons are in constant motion.

Electrons can be, in an instant, arranged in such a way they have a dipole. (Instantaneous dipole)

•The temporary dipole interacts with other temporary dipoles to cause attraction.

• exist between all molecules (polar and non polar)

•is the only attractive force between non polar atoms or molecules,

Temporary dipolesTemporary dipoles

Intermolecular ForcesIntermolecular Forces•Boiling point is dependant on the intermolecular forces

•Polar molecules have higher b.p. than nonpolar molecules.

Dipole - Dipole bondingDipole - Dipole bonding

Hydrogen bondingHydrogen bondingHydrogen bonding is an extreme form of dipole-

dipole attraction. Examples include: H2O, NH3, HF

Hydrogen Bonding :

the attractive force between hydrogen in a polar bond (particularly H-F, H-O, H-N bond) and an unshared electron pair on a nearby relatively small electronegative atom or ion

Very polar bond in H-F forms H bonds.The other hydrogen halides don’t form hydrogen bonds, since H-X bond is less polar. As well as that, their lone pairs are at higher energy levels. That makes the lone pairs bigger, and so they don't carry such an intensely concentrated negative charge for the hydrogens to be attracted to.

Hydrogen Bonding & Water

H-bonding causes the lower density of ice in comparison to liquid water, this means ice floats on water.

In most substances the molecules in the solid are more densely packed than in the liquid. A given mass of ice occupies a greater volume than that of liquid water.

This is because of an ordered open H-bonding arrangement in the solid (ice) in comparison to continual forming & breaking H-bonds as a liquid.

http://www.its.caltech.edu/~atomic/snowcrystals/photos/photos.htm

•Use your knowledge of bonding to compare the different boiling points of :Ionic/covalent network/metallic/and molecular substances

Test Info

electron configuration of atoms and ions of thefirst 36 elements (using s,p,d notation)

periodic trends in atomic radius, ionisationenergy and electronegativity

comparison of atomic and ionic radii

Lewis diagrams of molecules and ionsShape and polarity

Weaker Intermolecular Forces

Ion-Dipole Forces

An ion-dipole force is an attractive force that results from the electrostatic attraction between an ion and a neutral molecule that has a dipole. Most commonly found in solutions. Especially important for solutions of ionic compounds in polar liquids. A positive ion (cation) attracts the partially negative end of a neutral polar molecule. A negative ion (anion) attracts the partially positive end of a neutral polar molecule.

Ion-dipole attractions become stronger as either the charge on the ion increases, or as the magnitude of the dipole of the polar molecule increases.

Dipole-dipole Attractive Forces

A dipole-dipole force exists between neutral polar molecules Polar molecules attract one another when the partial positive charge on one

molecule is near the partial negative charge on the other molecule The polar molecules must be in close proximity for the dipole-dipole forces to

be significant Dipole-dipole forces are characteristically weaker than ion-dipole forces

Dipole-dipole forces increase with an increase in the polarity of the molecule

Boiling points increase for polar molecules of similar mass, but increasing dipole:

SubstanceMolecular

Mass (amu)Dipole

moment, u (D)Boiling Point

(°K)

Propane 44 0.1 231

Dimethyl ether

46 1.3 248

Methyl chloride

50 2.0 249

Acetaldehyde

44 2.7 294

Acetonitrile 41 3.9 355

London Dispersion Forces – significant only when molecules are close to each other

Prof. Fritz London

Due to electron repulsion, a temporary dipole on one atom can induce a similar dipole on a neighboring atom

Group 4A hydrides

Groups 4, 5, 6A hydrides

Van der Waals forces are made of dipole-dipole and London dispersion forces

Periodic Table TrendsPeriodic Table Trends

The Structure The Structure of Lithium of Lithium FluorideFluoride

Formal ChargeFormal Charge

Formal charge = (# valence e in neutral atom

# “assigned” valence e in molecule)

“assigned” elone pair e

bonded pair e2

1

VSEPR ModelVSEPR Model(“valence shell electron pair repulsion”)(“valence shell electron pair repulsion”)

Write Lewis structure Write Lewis structure Count Count ee pairs around atoms pairs around atoms

(single bonds, multiple bonds, and lone pairs (single bonds, multiple bonds, and lone pairs all count as 1 pair)all count as 1 pair)

Arrange bonds to minimize repulsion Arrange bonds to minimize repulsion between between ee pairs pairs

Lone pairs have greater repulsion than Lone pairs have greater repulsion than bonding pairsbonding pairs

Molecular Molecular Structure Structure of Methaneof Methane

The Molecular Structure of NHThe Molecular Structure of NH33

The Molecular Structure of HThe Molecular Structure of H22OO

The Bond Angles in the CHThe Bond Angles in the CH44, NH, NH33, and , and

HH22O MoleculesO Molecules

EElectron lectron PairsPairs

Bonding Pair

Lone pair

Possible Possible Electron Pair Electron Pair Arrangements Arrangements for XeFfor XeF44

Three Possible Arrangements of Three Possible Arrangements of the Electron Pairs in the Ithe Electron Pairs in the I33

-- Ion Ion

TThe Molecular he Molecular Structure of Structure of MethanolMethanol

Sizes of Sizes of Ions are Ions are Related Related to to Positions Positions of the of the Elements Elements in the in the Periodic Periodic TableTable

CompoundCompound Bond Length Bond Length ((Å)Å)

Electronegativity Electronegativity DifferenceDifference

Dipole Dipole Moment (D)Moment (D)

H-FH-F 0.920.92 1.91.9 1.821.82

H-ClH-Cl 1.271.27 0.90.9 1.081.08

H-BrH-Br 1.411.41 0.70.7 0.820.82

H-IH-I 1.611.61 0.40.4 0.440.44

Electronegativity difference decreases as bond length increases

Dipole Moment: µ = Qr

Dipole moment is defined as the magnitude of charge (Q) multiplied by the distance between the charges; units are D (Debye) = 3.36 x 1030 C.m

Prof. Peter DebyeNoble Prize 1936

When proton & electron 100 pm apart, the dipole moment is 4.80 D

4.8 D is a key reference value! It represents a pure charge of +1 and -1, which are 100 pm (100pm = 1Å) apart. The bond is said to be 100% ionic!

H-F; µ = 1.82 D (measured) bond length = 0.92 pmIf 100% ionic,µ = 92/100 (4.8 D) = 4.42 D% ionic = 1.82/4.42 x 100 = 41 % ionic

H-Cl; µ = 1.08 D (measured) bond length = 1.27 pmIf 100% ionic,µ = 127/100 (4.8 D) = 6.10 D% ionic = 1.08/6.10 x 100 = 18 % ionic

H-Br; µ = 0.82 D (measured) bond length = 1.41 pmIf 100% ionic,µ = 141/100 (4.8 D) = 6.77 D% ionic = 0.82/6.77 x 100 = 12 % ionic

Symmetrical molecules (e.g. CCl4, CH4) are non-polar. The four dipoles

are of equal magnitude and neutralize one another at the center of a

tetrahedron

Non-symmetrical molecules (e.g. CHCl3, CO(CH3)2, H2O) are Polar.

The dipoles are not all equal or in opposite directions (partial charges and

bond lengths are all different in C-Cl, C-H, C=O, C-H)

(H2O is a bent molecule not linear, see later notes)

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ATOMIC STRUCTURE electron configuration of atoms and ions of the