IB Chemistry ATOMIC THEORY

IB Chemistry

ATOMIC THEORY

Atomic Structure

Atomic StructureAtoms are very small ~ 10-10 metersAll atoms are made up of three sub-atomic particles: protons, neutrons and electrons

The protons and neutrons form a small positively charged nucleusThe electrons are in energy levels outside the nucleus

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

A meaningful way to consider the masses of the sub-atomic particles is to use relative masses

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.

No. of neutrons = Mass number – atomic number

So how can you work out the number of neutrons in an atom?

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?Example

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.

5. Which neutral element contains 11 electrons and 12 neutrons?

Atomic Structure - Questions

5. Copy and complete the following table:

Summary SlideAll atomic masses are relative to the mass of carbon-12.

Eg one hydrogen atom weighs 1/12 the mass of a carbon-12 atom.

IsotopesIsotopes 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 Cl3517

3717

17 protons17 electrons18 neutrons

17 protons17 electrons20 neutrons

The isotopes of chlorine are often referred to as chlorine-35 and chlorine-37

IsotopesIsotopes 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 Lower densities • lower melting and boiling points

Isotopes - Questions1. 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:a) one atom of carbon-12b) one atom of carbon-14c) one atom of uranium-235d) one atom of uranium-238

Mass SpectrometerThe 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 deflected by the magnetic field, and the amount of deflection depends upon the mass/charge ratio of the charged particle.

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)

or X (l) X (g)


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 the positive ions towards the magnetic field. The accelerated ions are focused and passed through a slit: this produces a narrow beam of ions.


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• the strength of the magnetic field is greater

Mass SpectrometerIf 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 sufficiently to reach the detector.

Mass SpectrometerDetection – ions that reach the detector cause electrons to be released in an ion-current detectorThe 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 spectrum.

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 session.

Atomic Structure – Mass Spectrometer

Isotopes of boron

m/z value 11 10Relative abundance %

18.7 81.3

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

= 205.7 + 813 100

= 1018.7 = 10.2 100

Mass Spectrometer – Questions

A mass spec chart for a sample of neon shows that it contains: 90.9% 20Ne 0.17% 21Ne 8.93% 22Ne

Calculate the relative atomic mass of neon

You must show all your working!


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

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

100

Ar= 20.18


Calculate the relative atomic mass of lead

You must show all your working!m/

e204 206 207 208

52.3

23.622.61.5


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)

100306 + 4861.6 + 4678.2 + 10878.4

100

20724.2100

Ar= 207.24

Energy LevelsElectrons 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 12 23 34 4

These sub-levels are assigned the letters, s, p, d, f

Energy LevelsEach type of sub-level can hold a different maximum number of electron.

Sub-levelMaximum number of electrons

s 2p 6d 10f 14

Energy LevelsThe energy of the sub-levels increases from s to p to d to f. The electrons fill up the lower energy sub-levels first.

Looking at this table can you work out in what order the electrons fill the sub-levels?

Energy LevelsLet’s take a look at the Periodic Table to see how this fits in.

Electronic StructureSo how do you write it?

1s2

Energy levelSub-level

Number of electrons

ExampleFor magnesium:1s2, 2s2, 2p6, 3s2

Electronic StructureThe 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 lower energy than the 3d sub-level.

Electronic StructureThe order in this the energy levels are filled is called the Aufbau Principle.Example (Sodium – 2, 8, 1)

Electronic Structure There are two exceptions to the Aufbau principle.The electronic structures of chromium and copper do not follow the pattern – they are anomalous.Chromium – 1s2, 2s2, 2p6, 3s2, 3p6, 3d5, 4s1

Copper – 1s2, 2s2, 2p6, 3s2. 3p6, 3d10, 4s1Write the electronic configuration for the following elements:a) hydrogen c) oxygen e) copperb) carbon d) aluminium f) fluorine

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

O atom O- ion

1s2 2s2 2p6

1s2 2s2 2p5

Negative ions: formed by the gain of e-

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

Chromium (Cr) - loss of 3 e-

Electronic Structure - Questions

Give the full electronic structure of the following positve ions:a) Mg2+ b) Ca2+ c) Al3+

Give the full electronic structure of the negative ions:a) Cl- b) Br- c) P3-

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 10S2- 16 16Sc3+ 21 45Ni2+ 30 26

OrbitalsThe 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 2p 3 6d 5 10f 7 14

OrbitalsThe orbitals in different sub-levels have different shapes:

• s orbitals1s 2s

• p orbitals

OrbitalsWithin 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

OrbitalsThe 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 between the paired electrons.

Electronic structure of carbon, 1s2, 2s2, 2p21s

2s2p

OrbitalsUsing boxes to represent orbitals, give the full electronic structure of the following atoms:a) lithium b) fluorine c)

potassiumd) nitrogen e) oxygen

1s2s

2p



Electronic structure of lithium: 1s2, 2s1

1s2s

2p



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

1s2s

2p



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

1s2s

2p3s

3p

4s



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

1s2s

2p



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

1s2s

2p

Ionization EnergyIonization 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 > 0Since all ionizations requires energy, they are endothermic processes and have a positive enthalpy change (ΔH) value.

Ionization EnergyThe value of the first ionization energy depends upon two main factors:The size of the nuclear chargeThe energy of the electron that has been removed (this depends upon its distance from the nucleus)

Ionization EnergyAs the size of the nuclear charge increases the force of the attraction between the negatively charged electrons and the positively charged nucleus increases.

+ +Small

nuclear charge

Large nuclear charge

Small

force of attraction

Smaller ionization

energy

Large force of

attractionGreater

ionization energy

Ionization energyAs 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 ionizati

on energy

Electrons further away from positive

nucleusSmall force of

attraction

Smaller ionizati

on energy

+

Ionization energy - Questions

Write an equation to represent the first ionization of:

a) aluminiumb) lithiumc) sodium

Trends across a PeriodGoing across a period, the size of the 1st 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.

+ + + +

Going across a Period

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, 2s2, 2p6, 3s1 494Mg 1s2, 2s2, 2p6, 3s2 736Al 1s2, 2s2, 2p6, 3s2, 3p1 577Si 1s2, 2s2, 2p6, 3s2, 3p2 786P 1s2, 2s2, 2p6, 3s2, 3p3 1060S 1s2, 2s2, 2p6, 3s2, 3p4 1000Cl 1s2, 2s2, 2p6, 3s2, 3p5 1260Ar 1s2, 2s2, 2p6, 3s2, 3p6 1520

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 remove an electron, hence a reduction in the ionisation energy.

3s3p

Phosphorus

3s3p

Sulphur

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.

Clue: which sub-level (s, p, d or f is the outer electron in?

Electronic structure

Ionisation energy/kJ mol-1

Na 1s2, 2s2, 2p6, 3s1 494Mg 1s2, 2s2, 2p6, 3s2 736

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

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

P 1s2, 2s2, 2p6, 3s2, 3p3 1060

S 1s2, 2s2, 2p6, 3s2, 3p4 1000

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

Ar 1s2, 2s2, 2p6, 3s2, 3p6 1520

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.

0500

10001500200025003000

0 5 10 15 20 25

Atomic number (Z)

Firs

t ion

isat

ion

ener

gy/k

J m

ol-1

H

He

Li

Be

B

CN

O

F

Ne

Na

Mg

AlSi

P

S

ClAr

K

Ca

Trends down a Group+

+

+

+

Down the Group

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 electrons.

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?

4. Complete Tasks

Successive Ionization energyDefinition: 2nd i.e.The energy per mole for the process X+

(g) X2+(g) +e-

And so on for further successive ionisation energies

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 removed form a new sublevel closer to the nucleus

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 closer to nucleus

Electron Affinity Energy Change per mole for:

X (g) + e-

X-(g)

That is, for the gaseous atoms to gain an electron to form anions

Electron Affinity

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

The second e.a is positive (endothermic) because an electron is being added to an ion which is already negative : repulsion occurs

IB Chemistry ATOMIC THEORY

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