Atomic Theory&Isotopes

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IB Chemistry

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Atomic Structure

Atoms are very small ~ 10

-10

meters

All atoms are made up of three sub-atomic

particles:

protons, neutrons and electrons

_{The protons and neutrons form a small}

positively charged nucleus

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Atomic Structure

The actual values of the masses and charges of the sub-atomic particles are shown below:

_{A meaningful way to consider the masses of the}

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Atomic Structure -

Definitions

Atomic number

(

Z

) is the number of protons

in the nucleus of an atom. The number of

protons equals the number of electrons in a

neutral atom

 _{N.B. No. of protons always equals the no. of} electrons in any neutral atom of an element.

Mass number

(

A

) is the sum of the number of

protons and the number of neutrons in the

nucleus of an atom.

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Atomic Structure -

Example

No. of neutrons = Mass number – atomic number

No. of neutron = Mass No. – Atomic No. = 23 – 11

= 12

So how can you work out the number of neutrons in

an atom?

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Atomic Structure -

Questions

1.

What are the three sub atomic particles that make up the atom?

2.

Draw a representation of the atom and labelling the sub-atomic particles.

3.

Draw a table to show the relative masses and charges of the sub-atomic particles.

4.

State the atomic number, mass number and

number of neutrons of: a) carbon, b) oxygen and c) selenium.

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Atomic Structure -

Questions

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Summary Slide

All atomic masses are relative to the

mass of carbon-12.

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Isotopes

Isotopes are atoms of the same element with the same atomic number, but different mass numbers, i.e. they have different numbers of neutrons.

Each atom of chlorine contains the following:

Cl Cl 35 17 37 17 17 protons 17 electrons 18 neutrons 17 protons 17 electrons 20 neutrons

The isotopes of chlorine are often referred to as

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Isotopes

Isotopes of an element have the same chemical properties because they have the same number of electrons. When a chemical reaction takes place, it is the electrons that are involved in the reactions. However isotopes of an element have the slightly different physical properties because they have different numbers of neutrons, hence different masses.

The isotopes of an element with fewer neutrons will have:

 _{Lower masses} _{• faster rate of diffusion}

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Isotopes -

Questions

1.

Explain what isotopes using hydrogen as an example.

2.

One isotope of the element chlorine, contains 20 neutrons. Which other element also contains 20 neutrons?

3.

State the number of protons, electrons and neutrons in:

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Mass Spectrometer

The mass spectrometer is an instrument used:  _{To measure the relative masses of isotopes}

 _{To find the relative abundance of the isotopes in a}

sample of an element

When charged particles pass through a magnetic field, the particles are

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Mass Spectrometer –

5 Stages

Once the sample of an element has been

placed in the mass spectrometer, it

undergoes five stages.

Vaporisation

– the sample has to be in

gaseous form. If the sample is a solid or

liquid, a heater is used to vaporise some of

the sample.

X

_(s)



X

_(g)

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Mass Spectrometer –

5 Stages

Ionization

– sample is bombarded by a

stream of high-energy electrons from

an electron gun, which ‘knock’ an

electron from an atom. This produces a

positive ion:

X

_(g)



X

+

(g)

+ e

-Acceleration – an electric field is used to accelerate

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Mass Spectrometer –

5 Stages

Deflection

–

The accelerated ions are deflected into the magnetic field. The amount of

deflection is greater when:

• the mass of the positive ion is less

• the charge on the positive ion is greater

• the velocity of the positive ion is less

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Mass Spectrometer

If all the ions are travelling at the same

velocity and carry the same charge, the

amount of deflection in a given magnetic field

depends upon the mass of the ion.

For a given magnetic field, only ions with a

particular

relative

mass (

m

) to charge (

z

)

ration – the m/z value – are deflected

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Mass Spectrometer

Detection

– ions that reach the detector

cause electrons to be released in an

ion-current detector

The number of electrons released, hence the

current produced is proportional to the

number of ions striking the detector.

The detector is linked to an amplifier and

then to a recorder: this converts the current

into a

peak

which is shown in the

mass

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Atomic Structure –

Mass Spectrometer

Name the five stages which the sample

undergoes in the mass spectrometer

and make brief notes of what you

remember under each stage.

Complete

Exercise 4, 5

and

6 in the

handbook. Any incomplete work to be

completed and handed in for next

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Atomic Structure –

Mass

Spectrometer



Isotopes of boron

m/z value 11 10

Relative

abundance % 18.7 81.3

A_r of boron = (11 x 18.7) + (10 x 81.3) (18.7 + 81.3)

= 205.7 + 813 100

= 1018.7 = 10.2

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Mass Spectrometer –

Questions

A mass spec chart for a sample of neon

Ne

Calculate the relative atomic mass of neon

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Mass Spectrometer –

Questions

 90.9% 20Ne  0.17% 21Ne  8.93% 22Ne

(90.9 x 20) + (0.17 x 21) + (8.93 x 22) 100

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Mass Spectrometer –

Questions

Calculate the

relative atomic

mass of lead

You must show all

your working!

m/e

204 206 207 208 52.3

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Mass Spectrometer –

Questions

 1.5% 204Pb  23.6% 206Pb  22.6% 207Pb  52.3% 208Pb

(1.5 x 204) + (23.6 x 206) + (22.6 x 207)+(52.3 x 208)

100 306 + 4861.6 + 4678.2 + 10878.4

100 20724.2100

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Energy Levels

Electrons go in shells or energy levels.

The energy levels are called

principle

energy levels

, 1 to 4.

The energy levels contain sub-levels.

Principle energy level

Number of sub-levels

1 1

2 2

3 3

4 4

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Energy Levels

Each type of sub-level can hold a

different maximum number of electron.

Sub-level

Maximum number of

electrons

s 2

p 6

d 10

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Energy Levels

The energy of the sub-levels increases

from

s

to

p

to

d

to

f

. The electrons fill

up the lower energy sub-levels first.

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Energy Levels

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Electronic Structure

So how do you write it?

1s

2

Energy level

Sub-level Number of _electrons

Example

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Electronic Structure

The electronic structure follows a pattern – the order of filling the sub-levels is 1s, 2s, 2p, 3s, 3p…

After this there is a break in the pattern, as that the 4s fills before 3d.

Taking a look at the table below can you work out why this is?

• This is because the 4s sub-level is of

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Electronic Structure

The order in this the energy levels are

filled is called the

Aufbau Principle

.

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Electronic Structure

There are two exceptions to the Aufbau

principle.

The electronic structures of chromium and

Write the electronic configuration for the following elements: a) hydrogen c) oxygen e) copper

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Electronic Structure

– of ions

When an atom loses or gains electrons

to form an ion, the electronic structure

changes:



Positive ions: formed by the loss of e

-1s2 2s2 2p6 3s1 

1s2 2s2 2p4 

Na atom Na+ ion

1s2 2s2 2p6

1s2 2s2 2p5

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-Electronic Structure

– of transition metals

With the transition metals it is the

4s

electrons that are lost first when they

form ions:



Titanium (Ti) - loss of 2 e

-1s2 2s2 2p6 3s2 3p6 3d2 4s2 

Ti atom Ti2+ ion

Cr atom Cr3+ ion

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

1s2 2s2 2p6 3s2 3p6 3d5 4s1  1s2 2s2 2p6 3s2 3p6 3d3

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-Electronic Structure

- Questions

Give the full electronic structure of the

following positve ions:

a) Mg

2+

b) Ca

2+

c) Al

3+

Give the full electronic structure of the

negative ions:

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3-Electronic Structure

- Questions

Copy and complete the following table:

Atomic

no. Mass no. No. of protons No. of neutrons No. of electrons Electronic structure Mg 12 1s2 2s2 2p6 3s2

Al3+ 27 10

S2- 16 16

Sc3+ 21 45

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Orbitals

The energy sub levels are made up of

orbitals, each which can hold a maximum of 2

electrons.

Different sub-levels have different number of

orbitals:

Sub-level No. of _orbitals Max. no. of _electrons

s 1 2

p 3 6

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Orbitals

The orbitals in different sub-levels have

different shapes:

• s orbitals

1s 2s

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Orbitals

Within a sub-level, the electrons occupy

orbitals as unpaired electrons rather than

paired electrons. (This is known as Hund’s

Rule).

We use boxes to represent orbitals:

1s

2s

2p

Electronic structure of carbon, 1s2, 2s2, 2p2

 

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Orbitals

The arrows represent the electrons in the

orbitals.

The direction of arrows indiactes the spin of

the electron.

Paired electrons will have opposite spin, as

this reduces the

mutual repulsion

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Orbitals

Using boxes to represent orbitals, give the full

electronic structure of the following atoms:

Electronic structure of fluorine: 1s2, 2s2, 2p5

1s

2s

2p

 

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Orbitals

Using boxes to represent orbitals, give the full

electronic structure of the following atoms:

Electronic structure of potassium: 1s2, 2s2, 2p6, 3s2, 3p6, 4s1

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Orbitals

Using boxes to represent orbitals, give the full

electronic structure of the following atoms:

Electronic structure of nitrogen: 1s2, 2s2, 2p3

1s

2s

2p

 

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Orbitals

Using boxes to represent orbitals, give the full

electronic structure of the following atoms:

Electronic structure of oxygen: 1s2, 2s2, 2p4

1s

2s

2p

 

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Ionization Energy

Ionization of an atom involves the loss of an electron to form a positive ion.

The first ionization energy is defined as the energy required to remove one mole of electrons from one mole of atoms of a gaseous element.

The first ionization energy of an atom can be represented by the following general equation:

X_(g)  X+ + e- ΔH > 0

Since all ionizations requires energy, they are

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Ionization Energy

The value of the first ionization energy

depends upon two main factors:

The size of the nuclear charge

The energy of the electron that has

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Ionization Energy

As the size of the nuclear charge increases the force of the attraction between the negatively charged

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Ionization energy

As the energy of the electron increases, the electron is farther away from the nucleus. As a result the

force of attraction between the nucleus and the electron decreases.

+ Electrons closer to

positive nucleus



Large force of attraction  Greater ionization energy Electrons further away from positive nucleus 

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Ionization energy -

Questions

Write an equation to represent the first

ionization of:

a) aluminium

b) lithium

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Trends across a Period

Going across a period, the size of the 1

st

ionisation energy shows a general

increase

.

This is because the electron comes from the

same energy level, but the size of the nuclear

charge increases.

+ +

+

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Trends across a Period

(2 exceptions)

The first ionisation of Al is less than that of Mg, despite the increase in the nuclear charge.

The reason for this is that the outer electron

removed from Al is in a higher sub-level: the electron removed from Al is a 3p electron, whereas that

removed from Mg is a 3s.

Electronic structure Ionisation energy/kJ mol-1

Na 1s2_{, 2s}2_{, 2p}6_{, 3s}1 ₄₉₄

Mg 1s2_{, 2s}2_{, 2p}6_{, 3s}2 ₇₃₆

Al 1s2_{, 2s}2_{, 2p}6_{, 3s}2_{, 3p}1 ₅₇₇

Si 1s2, 2s2, 2p6, 3s2, 3p2 786

P 1s2_{, 2s}2_{, 2p}6_{, 3s}2_{, 3p}3 ₁₀₆₀

S 1s2_{, 2s}2_{, 2p}6_{, 3s}2_{, 3p}4 ₁₀₀₀

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Trends across a Period

(2 exceptions) The first ionisation energy of S is less than that of P, despite the increase in the nuclear charge.

In both cases the electron removed is from the 3p sub-level. However the 3p electron removed from S is a

paired electron, whereas the 3p electron removed from P

is an unpaired electron.

When the electrons are paired the extra mutual repulsion results in less energy being required to

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Trends across a Period -

Questions

There is a break in this general trend going across a Period.

Look at the table below and point out where the break in the the trend is and try to give an explanation.

Electronic structure Ionisation energy/kJ _mol_-1 Na 1s2_{, 2s}2_{, 2p}6_{, 3s}1 ₄₉₄

Mg 1s2_{, 2s}2_{, 2p}6_{, 3s}2 ₇₃₆

Al 1s2, 2s2, 2p6, 3s2, 3p1 577

Si 1s2, 2s2, 2p6, 3s2, 3p2 786

P 1s2_{, 2s}2_{, 2p}6_{, 3s}2_{, 3p}3₁₀₆₀

S 1s2_{, 2s}2_{, 2p}6_{, 3s}2_{, 3p}4₁₀₀₀

Cl 1s2, 2s2, 2p6, 3s2, 3p5 1260

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Trends across a Period -

Questions

Now take a look at the graph below:

a) Explain what the graph shows in as much detail as possible

b) There is one other break in the general pattern going across a Period. What is it and explain why that is.

0 500 1000 1500 2000 2500 3000

0 5 10 15 20 25

Atomic number (Z)

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Trends down a Group

+ + + + D ow

n t_he

G

ro

up

Ionization energy

decreases

going

down a Group.

Going down a Group in the Periodic

Table, the electron removed during

the first ionization is from a higher

energy level and hence it is further

from the nucleus.

The nuclear charge also increases,

but the effect of the increased

nuclear charge is reduced by the

inner electrons which

shield

the outer

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Ionization energy

- Questions

1. Explain why sodium has a higher first

ionization energy than potassium.

2. Explain why the first ionization energy

of boron is less than that of beryllium.

3. Why does helium have the highest

first ionisation energy of all the

elements?

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Successive Ionization energy

Definition: 2

nd

i.e.

The energy per mole for the process

X

+

(g)

X

2+(g)

+e

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Successive Ionization energy



Successive i.e’s increases because

electrons are being removed from

increasingly positive ions.



Therefore, nuclear attraction is

greater.



Large jumps seen when electron is

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Successive Ionization energy

Succesive ionisation energies of Calcium

0 1 2 3 4 5 6

0 2 4 6 8 10 12 14 16 18 20

Number of electron removed

2nd i.e higher than first – electron has greater pull from nucleus

Large increase between 4th

and 3rd shells – electron

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Electron Affinity



Energy Change per mole for:

X

_(g)

+ e

-

X

-(g)

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Electron Affinity

The first e.a is negative (exothermic) because the electron is attracted to the positive charge on the atom’s nucleus.