Atomic Theory

Unit 7

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Unit 7, 2

The Atom

An atom is the smallest part of any element that can take part in a chemical reaction.

An atom is the smallest particle of an element which still retains the properties of that element.

Atoms are extremely small.

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Unit 7, 3

Structure of the Atom

24Mass is measured in atomic mass unit 1 a.m.u. 1.660 10 kg.

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Unit 7, 4

Elements and Compounds

An element is a substance that cannot be split into a simpler substance by chemical means.

A compound is a substance which is made up of two or more different elements combined together chemically.

Most elements from the periodic table are found in nature as compounds very few are found in elemental form.

Element

Compounds

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Unit 7, 5

Molecules

A molecule is a group of atoms joined together.

It is the smallest part of an element or compound that can exist independently.

There are two types of molecules: Molecules of elements:

(All the same atoms in the molecule)e.g. Cl2, S8.

Molecules of compounds:(Different atoms in the molecule)e.g. H2O, HCl, H2SO4.

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Unit 7, 6

Diatomic Elements

Elements that end in –ine and –gen exist as diatomic molecules in their natural form.

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Unit 7, 7

Atoms, Molecules, Elements & Compounds

Atoms

Molecules

Compounds

Elements

of

of

of

(Only 1)

(2 or more)

(all the same)

(different)

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Unit 7, 8

Summary

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Unit 7, 9

The Periodic Table

The Periodic table is the chemist’s compass to a vast array of compounds.

In the modern periodic table the elements are listed in order of increasing atomic number, and arranged in order of the numbers of electrons in their outer main energy level.

These are referred to as the group number and the period number.

The periodic table is divided into two main sections metals and non – metals. A memory aid to where this division occurs is :Being Silly Assures Teachers Attention.

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Unit 7, 10

The Periodic TableP.T.E

Unit 7, 10

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Unit 7, 11

Some of the Named GroupsThe Named GroupsP.T.E

Unit 7, 11

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Unit 7, 12

Group Number

The group number of an element is equal to the numbers of electrons in the outer main energy level (shell) of an atom of that element.

The named groups are:Group one – Alkali metal.Group two – Alkaline earth metals.Group seven – Halogens. Group eight – Noble gases.

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Unit 7, 13

Period

The period number of an element is the number of main energy levels (shells) in an atom.

Period 1

Period 2

Period 3

Period 4

Period 7

Period 6

Period 5

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Unit 7, 14

Solids, Liquids and Gases

H He

Li Be B C N O F Ne

NaMg

Al Si P S Cl Ar

K Ca Sc Ti V CrMn

Fe Co Ni Cu Zn Ga Ge As Se Br Kr

Rb Sr Y Zr NbMo

Tc Ru Rh Pd Ag Cd In Sn Sb Te I Xe

Cs Ba La Hf Ta W Re Os Ir Pt Au Hg Tl Pb Bi Po At Rn

Fr Ra Ac RfDb

Sg Bh Hs Mt

Solids, Liquids and Gases

SolidsLiquids

Gases

Unit 7, 14

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Unit 7, 15

Ions

Ions are atoms or groups of atoms which have lost or gained electrons and hence have a charge.

A positive ion (cation) is formed when an atom(s) loses electrons, they are denoted by the following notation Xn+.

A negative ion (anion) is formed when an atom(s) gains electrons, they are denoted by the following notation Xn-.

Simple ions (ions of 1 atom only) can be worked out from their valency (and from the periodic table).

The valency of an element is equal to the number of electrons which an atom must lose or gain in order to attain noble gas configuration.

Simple Ions Simple IonsP.T.E

Hydrogen forms the H ion except when with a metal it becomes the hydride (H ) ion.

All form postive ions.

metalsUnit 7, 16

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Unit 7, 17

Transition Metal Ions

3

2

Transition metals can have more than one valency, so if this is the case the

valency is written in roman numerals after the metal.

FeCl Iron (III) chloride

FeCl Iron(II) chloride

Note:

If the valency is one the roman numerals are not required.

If there is only one valency for the metal, roman numerals are not

required either.

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Unit 7, 18

Polyatomic Ions

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Unit 7, 19

Making Formulas for Compounds

Get the correct ions from: Periodic Table of Elements. Transition metals table. Polyatomic ions table.

Put the positive ion first. If the charges are different cross over the

magnitude (number part) of the charges. If you require more than one of a

polyatomic ion make sure to put it in brackets before adding the subscript.

There should be no charges on a compound.

2 2Mg O

MgO

2+

2-

What is the formula for Magnesium oxide?

Write down the ions.

Mg got from its position in the periodic table. (Group two)

O got from its position in the periodic table. (Group six)

The

Problem:

Solution:

charges are equal so simply write down the ions

beside each other.

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Unit 7, 20

3 2

2 3

Al O

Al O

3+

2

What is the formula for aluminium oxide?

Write down the ions.

Al got from its position in the periodic table (Group 3)

O got from its position in the periodic table (Group 6)

The charge

Problem:

Solution:

s are not equal so cross the magnitude of the

charges over.

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Unit 7, 21

Na OH

NaOH

+

What is the formula for sodium hydroxide?

Write down the ions.

Na got from its position in the periodic table (Group 3)

OH got from the table polyatomic ions (Learn!)

The charges are equ

Problem:

Solution:

al so simply write the ions down beside

each other.

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Unit 7, 22

2+

34

What is the formula for Magnesium phosphate?

Write down the ions.

Mg got from its position in the periodic table. (Group 2)

PO got from the table of polyatomic ions. (Learn!)

The charg

Problem:

Solution:

es are not equal so cross the magnitude of the

charges over.

2 34

3 4 2

Mg PO

Mg (PO )

N.B. If you require more than one of a polyatomic ion you must put it in brackets.

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Unit 7, 23

3+

What is the formula for Iron (III) chloride.

Write down the ions.

Fe the three in roman numerals gives the charge of a

transition metal with variable valency

Cl got from its position in

Problem:

Solution:

the periodic table. (Group 7)

The charges are not equal so cross the magnitude of the

charges over.

3

3

Fe Cl

FeCl

N.B. When we require only one of an item there is no need to fill in a one, it is assumed

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Unit 7, 24

+

What is the formula for Potassium hydride.

Write down the ions.

K got from its postion in the periodic table. (Group one)

H always H when joined to a metal.

The charges are equal so just

Problem:

Solution:

write them down beside each other.

K H

KH

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Unit 7, 25

Try Some!Try Some!

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Unit 7, 26

Naming Compounds Summary

Type II metals are the

transition metals.

The prefixes are:

mon one

di two

tri three

tetra four

pent five

Notes:

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Unit 7, 27

Atomic Number

The modern periodic table has the elements arranged in order of increasing atomic number. (Smaller of the two)The atomic number of an element in the Periodic Table tells us the number of protons in the nucleus.

i.e. sodium is the eleventh element in the periodic table and therefore it has 11 protons in its nucleus.

008_PERIODICPROP.MOV

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Unit 7, 28

Atomic Mass Number

The atomic mass number of an element is the sum of the numbers of protons and neutrons in the nucleus of an atom.

For example, sodium has 11 protons and 12 neutrons in the nucleus. Therefore its mass number is 23 or we say that the mass of the atom is 23 atomic mass units.

A special unit called the atomic mass unit is used as the masses of the atoms are so small.

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Unit 7, 29

Nuclear Formula

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Unit 7, 30

Nuclear Formula

It is called a nuclear formula as it only gives information about the nucleus of the atom or ion, you have to deduce the information for electrons.

Remember: All atoms are neutral. Positive ions have lost electrons. Negative ions have gained electrons.

2

1123

1939

612

How many protons neutrons and electrons are in each of the

following:

Na; K ; C; O .

Na 11 protons; 12 neutrons and 11 electrons.

K 19 protons; 20 neutrons and 18 electrons.

C

Problem:

Solution:

8 216

6 protons; 6 neutrons and 6 electrons.

O 8 protons; 8 neutrons and 10 electrons.

Unit 7, 31

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Unit 7, 32

Isotopes

In 1919 an English chemist called Francis William Aston built an instrument called a mass spectrometer to measure the masses of atoms.

He started work with a sample of neon gas and discovered something very unusual. He found that that neon gas consisted of two varieties of neon atom.

One type of neon atom had a mass number of 20 and the other type had a mass number of 22.

He concluded that neon gas consisted of atoms of neon that differed in the number of neutrons in the nucleus.

20 2210 10These two varieties Ne and Ne are said to be

of neon.isotopes

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Unit 7, 33

Isotopes of Chlorine & Carbon

35 3717 17

Having identified two isotopes of neon, Aston also identified two

isotopes of chlorine: Cl and Cl, commonly called chlorine 35

and chlorine 37.

12 13 146 6 6

There are many other examples of isotopes, e.g. carbon

has three isotopes: C, C and C., i.e. carbon 12,

carbon 13 and carbon 14, respectively.

are atoms of the same

element (i.e. they have the same

atomic number) that have different

mass numbers due to the different

number of neutrons in the nucleus.

Isotopes

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Unit 7, 34

Isotopes of Lithium

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Unit 7, 35

Isotopes of Hydrogen

Three isotopes have been found and their nuclei are shown above.

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Unit 7, 36

% of Each Isotope Present

For his work on the discovery of the isotopes Aston received the 1922 Nobel Prize in chemistry.

Not only did Aston detect the presence of isotopes, but he also determined the percentages of each of the isotopes present.

For example in his study of chlorine gas he found that there were approximately three times as many chlorine – 35 atoms as there were chlorine – 37 atoms. He was then able to work out the average mass of an atom of chlorine.

The periodic table contains this figure for the average mass of a atom of the element concerned. This is why the values given in the periodic table are all decimals.

His method of calculation is shown in the next example.

35 3717 17A sample of chlorine is found to consist of 75% Cl and 25% Cl .

Calculate the average mass of an atom of chlorine.

In 100% of chlorine there are:

75% of atoms with a mass of 35 75

Example:

Solution:

% 35 26.25

25% of atoms with a mass of 37 25% 37 9.25

100% 35.5

Average mass of chlorine 35.5 a.m.u.

Unit 7, 37

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Unit 7, 38

Relative Atomic Mass

If you locate chlorine in the periodic table you will find that 35.45 is the number given under the symbol for chlorine.

This ‘average mass’ of an atom is measured relative to the mass of the carbon – 12 isotope.

For this reason it is called the relative atomic mass. The symbol for relative atomic mass is Ar.

Since the relative atomic mass is the ratio of two masses, it has no units. Therefore, we say that the average mass of an atom of chlorine is 35.5 a.m.u but its relative atomic mass is 35.5.

The relative atomic mass of an element is the mass of an atom of that

element compared to one twelved the mass of the carbon 12 isotope.

69 7131 31

6931

The element gallium (Ga) consists of 60% Ga and 40% Ga.

State the number of neutrons in each of these isotopes.

Calculate the relative atomic mass of gallium.

Ga co

Problem:

(i)

(ii)

Solution:

(i)7131

6931

7131

ntains 38 neutrons. (69 31 38)

Ga contains 40 neutrons. (71 31 40)

60% of Ga 60% 69 41.4

40% of Ga 40% 71 28.4

100% 69.8

No units for relative atomic mass

(ii)

N.B:

Unit 7, 39

Unit 7, 40

The Mass SpectrometerThe response of the ion detector (intensity of lines on photographic plate) is converted to a scale of relative numbers of atoms. (i.e. % of

each isotope present)

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Unit 7, 41

The percent natural abundances of the mercury isotopes are:

Mass spectrum for mercury.

80196

80198

80199

80200

80201

80202

80204

Hg, 0.146%;

Hg, 10.02%;

Hg, 16.84%;

Hg, 23.13%;

Hg, 13.22%;

Hg, 29.80%;

Hg, 6.85%.

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Unit 7, 42

Average Mass of Mercury

80196

80198

80199

80200

80201

Averaging the seven isotopes from the spectrum:

0.146% of Hg 0.146% 196 0.2861

10.02% of Hg 10.02% 198 19.8396

16.84% of Hg 16.84% 199 33.5116

23.13% of Hg 23.13% 200 46.2600

13.22% of

80202

80204

Hg 13.22% 201 26.5722

29.80% of Hg 29.80% 202 60.1960

6.85% of Hg 6.85% 204 13.974

100% 200.59

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Unit 7, 43

Calculating Molecular Mass

2

3

r

r

r

r

Add up each individual atom in the molecule remembering to

multiply it by the number to the right of

H O 1 16

CH COOH 12 1 12 16 16

it.

The symbol for molecular mass is M .

M of ( ) 18

M of ( ) 60

M

2

3 1

2

2r 4

of ( ) 71

M of (

Cl 35.5

H SO 1 32 16

2

) ) 982 4(

3

r 3

r

r 3

Calculate the percentage by mass of calcium in calcium carbonate [CaCO ].

[C = 12; O = 16; Ca = 40.]

M of CaCO 40 12 3(16) 100

M of Ca 40% of Ca 100 100 40%

M of CaCO 100

Problem 1:

Solution:

Problem 2 :

2 5

r 2 5

r

r 2 5

Calculate the percentage by mass of hydrogen in ethanol [C H OH].

[C = 12; O = 16; H = 1.]

M of C H OH 2(12) 5(1) 16 1 46

M of H 6(1)% of H 100 100 13.04%

M of C H OH 46

Solution:

Unit 7, 44

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Unit 7, 45

The Arrangement of Electrons

Up to this we said that electrons orbit the nucleus in shells (Bohr’s theory).

Each shell contains electrons with a certain amount of energy and electrons normally occupy the lowest available energy level.

The ground state of an atom refers to its state when all of its electrons are in their lowest available energy levels.

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Unit 7, 46

In this situation the first shell n = 1 could hold 2n2 = 2 electrons.

In this situation the second shell n = 2 could hold 2n2 = 8 electrons.

In this situation the third shell n = 3 could hold 2n2 = 18 electrons.

In this situation the fourth shell n = 4 could hold 2n2 = 32 electrons.

This works well to explain some topics in chemistry, but a more detailed arrangement is sometimes necessary.

No. of Electrons in each Shell

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Unit 7, 47

Bohr Model of the Atom

In a Bohr atom, the electron is a particle that travels in specific, fixed orbits, but never in the space between orbits.

This arrangement expresses energy quantization, and accounts for the atomic emission spectra.

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Unit 7, 48

A Ladder for Distance, Bohr Model for Energy

Bohr orbits are like steps in a ladder. It is possible to be on one step or another, but it is impossible to be between steps.

Unlike the ladder Bohr orbitals do not have equal spacing between the orbits.

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Unit 7, 49

Energy Absorption and Emission

When a hydrogen atom absorbs energy, an electron is excited to a higher energy level.

The electron is then in an unstable and temporary level.

The electron falls back to the lower energy level and emits a photon of light (or some other form of radiation).

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Unit 7, 50

Evidence for the Existence of Energy Levels

When atoms are excited (i.e. given energy) by heating them or subjecting them to electrical discharge, they usually emit light or some other form of radiation.

A spectrum of light from a bulb would produce a continuous spectrum from red through to the various colours to violet.

The spectrum from excited atoms are not continuous but consist of a number of distinct lines each corresponding to a definite frequency of light. (Each colour of light has a different frequency)

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Unit 7, 51

Spectrum of Ordinary White Light

A continuous spectrum

with no gaps, with one

colour (corresponding

to a particular frequency)

merging into the next.

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Unit 7, 52

The Emission Spectrum of Helium

Unlike the continuous spectrum

for white light there are gaps

with distinct lines of particular

frequencies present and other

frequencies obviously missing

(black parts).

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Unit 7, 53

Spectrum of Hydrogen

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Unit 7, 54

Line Spectrum of Selected Elements

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Unit 7, 55

Visible Lines in the Hydrogen Spectrum

Each Bohr orbit has a distinct, fixed energy. When an electron relaxes from a high energy orbit to a low energy orbit, light (or some other form of radiation) is released.

The 486 nm line corresponds to an electron falling from the n = 4 to the n = 2 orbit.

The 657 nm line corresponds to an electron relaxing from the n = 3 to n = 2 orbit.

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Unit 7, 56

Model of Hydrogen & Emission Spectrum

A portion of the hydrogen atom model is shown with the nucleus at the centre of the atom and with the electron in one of a set of discrete orbits.

When the atom is excited, the electron jumps to a higher orbit (black arrows).

Transitions in which the electron falls to the second level are accompanied by the emission of light.

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Unit 7, 57

Orbital Energy Levels for Hydrogen

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Unit 7, 58

Energy Levels and Spectral lines for hydrogen

Visible light given out when

electrons fall to level two.

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Unit 7, 59

Evidence for the existence of energy levels

Continuous Spectrum (bulb)

Emission Spectrum (Hydrogen)

The fact that the spectrum of hydrogen consists of distinct lines indicates that only certain energy emissions are possible.

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Unit 7, 60

This was accounted for as follows by Bohr

In an atom the electrons revolve around the nucleus only in certain allowed ‘orbits’ or shells.

While in a particular shell, an electron has a definite amount of energy.

Electrons normally occupy the lowest available ‘energy level’ or ground state.

When energy is given to an atom one or more electrons are promoted to a higher energy level or ‘excited state’, such a state is unstable and temporary.

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Unit 7, 61

When the electron falls back to the lower energy level the energy difference between the two levels is emitted as a unit of light (a ‘photon’) or some other form of radiation, such as infrared or ultra violet.

Absorption & Emission

1E 1E

2E 2E

hf

hf

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Unit 7, 62

Equation for Energy Difference

The emitted radiation has a definite frequency – corresponding to the energy difference between the two levels.

E2 – E1= hf

E2 – energy of the higher level.

E1 – energy of the lower level.

h – Planck's constant.f – frequency of the emitted radiation.

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Unit 7, 63

Setup to Examine Hydrogen Spectrum

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Unit 7, 64

Apparatus for Observing Spectra

Spectrometer

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Unit 7, 65

Energy Levels & Sublevels

Bohr’s theory explains the various lines in the hydrogen spectrum very well, but hydrogen is relatively simple, with only one electron, for the other elements a lot of modification was required.

A close study showed the main energy level (principle level) had to be split into one or more sublevels.

The number of sublevels is equal to the number of the main energy level.

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Unit 7, 66

Energy levels are split into sublevels as follows: n = 1 has 1 sublevel, called 1s. n = 2 has 2 sublevels called 2s and 2p. n = 3 has 3 sublevels called 3s, 3p and 3d. n = 4 has 4 sublevels called 4s, 4p, 4d and 4f. A s – sublevel holds 2 electrons. A p – sublevel holds 6 electrons. A d – sublevel holds 10 electrons. A f – sublevel holds 14 electrons.

An energy level is the fixed amount of energy an electron has due to its position in an atom.

Energy Levels

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Unit 7, 67

Shells are Organised into Subshells

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Unit 7, 68

Arrangement of Energy Levels

1s

2s

2p

3s

3p

3d4p

4s

4d

4f

n = 1

n = 2

n = 3

n = 4

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Unit 7, 69

Electronic Configurations

Arrangement of electrons in an atom or ion.

Memory Aid:

This memory aid is to ensure that the

is filled the .4s before 3d

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Unit 7, 70

Examples:

Li 1s2, 2s1

F 1s2, 2s2, 2p5

Cl 1s2, 2s2, 2p6, 3s2, 3p5

Ar 1s2, 2s2, 2p6, 3s2, 3p6

Two exceptions are: Cr 1s2, 2s2, 2p6, 3s2, 3p6, 4s1, 3d5

Instead of 1s2, 2s2, 2p6, 3s2, 3p6, 4s2, 3d4

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

Instead of 1s2, 2s2, 2p6, 3s2, 3p6, 4s2, 3d9

(Full and half-full sublevels have extra stability)

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Unit 7, 71

Valence Electrons

Note: that the valence electrons are in the highest main energy level. The valence electrons are the first ones to be removed.

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Unit 7, 72

Electronic Configurations of ions:

A positive ion has lost electronsA negative ion has gained electronsIons lose or gain electrons to

become more stable, so when the electronic configuration is written out - the part in square brackets should be a noble gas.

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Unit 7, 73

Examples

Mg2+ [1s2, 2s2, 2p6]2+

1s2, 2s2, 2p6 = Ne

Cl- [1s2, 2s2, 2p6, 3s2, 3p6]-

K+ [1s2, 2s2, 2p6, 3s2, 3p6]+

1s2, 2s2, 2p6, 3s2, 3p6 = Ar

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Unit 7, 74

Filling of Outer Levels

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Unit 7, 75

Blocks of Sublevels

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Unit 7, 76

d – block

From Sc to Zn(This is just the first period of them)

The d – block are so called as the change/build up in electron structure takes place in the d – orbitals.

The d – block contains the transition elements.

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Unit 7, 77

Transition Metals

Form coloured compounds Cr3+ Green; Cu2+ Blue

Have variable valencies Fe2+, Iron (II) ion; Fe3+ Iron (III) ion.

Can act as catalysts A catalyst speeds up or slows down a reaction and

does not itself take part in the reaction. Exceptions to these are Sc and Zn as they do not form

coloured compounds or exhibit variable valencies. Sometimes transition metals are said to go from

Ti to Cu ( excluding Sc and Zn)

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Unit 7, 78

Atomic orbitals

Regions in space where a electron is most likely to be found.

Each orbital can only hold a maximum of two electrons.

This they can only do if they have opposite spin (Hund’s rule).

Electrons fill orbitals of equal energy singly before filling them in pairs (Alfbau principle).

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Unit 7, 79

s - orbitals

s – orbitals are spherical in shape.

The 2s – orbital is similar in shape to the 1s orbital, but larger in size.

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Unit 7, 80

p - orbitals

Dumb-bell in shape

3 p – orbitals to make up a p sublevel.

px

py

pz

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Unit 7, 81

Electronic Configurations with Orbitals of Equal Energy

Remember Hund’s Rule and Aufbau Principle.

The following is the electron distribution of the first ten elements in the periodic table

A represents an orbital. The arrows represent the spin

direction.

043_ElectronConfig.mov

He

H

Ne

O

N

B

Be

Li

C

F

1s 2s 2px 2py 2pZ

Orbital Diagrams when

Filling Orbitals of Equal Energy

Unit 7, 82

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Unit 7, 83

Quantum Numbers:

Set of numbers used to describe each electron uniquely in an atom.

There are four in total:Principle Quantum Number.Secondary Quantum Number.Magnetic Quantum Number.Spin Quantum Number.

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Unit 7, 84

Principal Quantum Number:

The number of the main energy level that an electron is in.

Values are 1, 2, 3 and 4 corresponding to the main energy levels n =1, n = 2, n = 3 and n = 4 respectively.

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Unit 7, 85

Secondary Quantum Number

The number of the energy sublevel (that an electron is in).

Values 0, 1, 2 and 3 corresponding to the sublevels s, p, d and f, respectively.

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Unit 7, 86

Magnetic Quantum Number

The number of the orbital (that an electron is in).

An s sublevel, has one orbital, assigned a magnetic quantum number of 0.

A p sublevel has 3 orbitals, assigned magnetic quantum numbers of either –1, 0, +1, corresponding to px, py or pz respectively.

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Unit 7, 87

Spin Quantum Number

Gives the direction of the spin of an electron in an orbital.

The first electron in the orbital is assigned a value of +½ and the second -½.

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Unit 7, 88

Summary Illustration

1x2p

Principal

Secondary

Spin

Magnetic

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Unit 7, 89

Quantum Number Summary Table

Principal(main energy level)

Secondary

(sublevel)

Magnetic

(orbital)

Spin(direction of spin)

1 1s 0 1s 0 First +½

Second –½

2

2s 0 2s 0 First +½

Second –½

2p 1

2px –1 First +½

Second –½

2py 0 First +½

Second –½

2pz +1 First +½

Second –½

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Unit 7, 90

Assigning Quantum Numbers

37

2 1

The outer electron in Li:

The electronic configuration of Li is:

1s , 2s

Principal quantum no.: 2 (Second main energy level)

Secondary quantum no.: 0 (s sublevel)

Magnetic quantum no.: 0 (s orbital

full

)

Spin quantum number.: ½ (first electron in the orbital)

Li 2, 0, 0, ½

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Unit 7, 91

Assigning Quantum Numbers

1020

2 2 2 2 2x y z

The outer electron in Ne

The electronic configuration of Ne is:

1s , 2s , 2p , 2p , 2p

Principal quantum no.: 2 (Second main energy level)

Secondary quantum no.: 1 (p sublevel)

Magnetic quantu

full

zm no.: 1 (p orbital)

Spin quantum number.: ½ (second electron in the orbital)

Ne 2, 1, 1, ½.

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Unit 7, 92

Ionisation Energy

The first ionisation energy of an element is the minimum energy required to remove the most loosely bound electron from an isolated atom of that element in its gaseous state.

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Unit 7, 93

Graph of the first Twenty Ionisation Energies

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Unit 7, 94

Explaining the Graph1. The maximum values are for the noble gases.

Reason: Their atoms are very stable because of their electronic configuration [full outer (sub) level], so it is difficult to remove an electron.

2. The minimum values are for the group one metals (alkali metals).Reason: Their atoms have only one electron in their outer level, so it

is easily removed (as when this is lost it will have noble gas configuration.) This is why group one are so reactive.

3. In general, ionisation energies increase in moving across a period from the alkali metal to the next noble gas.Reason: 1. Increase in nuclear charge.

(greater pull for electrons)2. Decrease in atomic radius.

4. Ionisation energies gradually decrease in moving down a group.Reason: 1. Increase in atomic radius.

2. Screening effect.(This is where the inner shell or shells of electrons help

to shield the outer electrons from the positive charge in the nucleus.

P.T.E

Unit 7, 95

Exceptions to Rule to 3 – Across a Period

There are two exceptions to this generalisation:(a) Group two elements (e.g. Be, Mg) have abnormally high values. This is because the most loosely bound electron comes from a full s orbital.

(e.g. 1s2, 2s2, 2p6, 3s2 in Mg) which is a relatively stable state. When the next element in each case (B, Al) is be ionised, the electron being removed

is the single electron in the p – orbital (e.g. 1s2, 2s2, 2p6, 3s2, 3p1 in Al).

(b) Group five elements also show abnormally high values (e.g. N and P). The reason here is that the electrons being removed are from exactly half – filled p –orbitals, (e.g. 1s2, 2s2, 2p6, 3s2, 3p3 in P)

and the half filled orbitals are the next most stable state after that of completely filled orbitals.

P.T.E

Unit 7, 96

Ionisation Energy Trends - Summary

Increase going across a period. Increase in nuclear charge. Decrease in atomic radius.

Decrease going down a group. Increase in atomic radius. Screening effect.

Exceptions, Group 2. Full (outer) sublevel.

Exceptions, Group 5. Half full (outer) sublevel.

2 2

2 2 6 2

2 2 6 2 6 2

Be 1s 2sMg 1s 2s 2p 3sCa 1s 2s 2p 3s 3p 4s

2 2 3

2 2 6 2 3

N 1s 2s 2pP 1s 2s 2p 3s 3p

Ionisation Energy TrendsP.T.E

Unit 7, 97

Trends in Ionisation Energies

Atomic RadiiP.T.E

Unit 7, 98

Atomic Radii

P.T.E

Unit 7, 99

Atomic Radii Versus Ionic Radii

1The following table gives the first ionisation energies , in KJ mol ,of

the elements in the second period of the Periodic table.

Li Be B C N O F Ne

519 900 799 1090 1400 1310 1680 2080

Explain the factors

Example 1:

(i) which account for the trend in ionisation energies

across a period.

Explain why the values for boron and oxygen are exceptional.

Increase in nuclear charge.

Decrease in atomic radius.

(ii)

Solution:

(i)

(ii) The values for boron and oxygen seem exceptional as the values

for the atoms before them have abnormal values, Be due to the fact

the electron is being removed from a full s - sublevel and N as the

electron is being removed from a half full p - sublevel, both of which have

extra stability.Unit 7, 100

Explain why the first ionisation energy of oxygen atoms is greater that

that of chlorine atoms.

Chlorine is below Oxygen in the periodic table and ionisation energies

decrease as you

Example 2 :

Solution:

1

1

go downwards due to:

1. Increase in atomic radius

2. Screening effect

The first ionisation energy of Sodium is 496 KJ mol , and the second ionsiation

energy is 4562 KJ mol . Account for the la

Example 3:

st 2 2 6 2 2 6

nd 2 2 2 2

1

2 56

rge difference.

1 ionisation energy: Na(1s ,2s ,2p , ) Na (1s ,2s ,2p ) e

2 ionisation energy: Na (1s ,2s , ) Na (1s ,2s ,2p ) e

The first ionisation energy is removing an electron

3

f

p

r

2

om

s

Solution:

a 3s orbital after

which the electronic configuration will be that of a noble gas, with a full

outer level.

The second ionisation energy is removing an electron from a full p - orbital

in a full level which is much closer to the nucleus hence a much higher value.Unit 7, 101

P.T.E

Unit 7, 102

Higher Ionisation Energy Levels for the Third

Period