atomic structure chemistry a level

8/12/2019 Atomic Structure CHEMISTRY A LEVEL

http://slidepdf.com/reader/full/atomic-structure-chemistry-a-level 1/48

Atomic Structure - Questions

1. What are the three sub atomic particles that makeup 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 andc) selenium.

5. Which neutral element contains 11 electrons and12 neutrons?



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

Each atom of chlorinecontains the following:

Cl Cl35

17

37

17

17 protons17 electrons

18 neutrons

17 protons17 electrons

20 neutrons

The isotopes of chlorine are often referred to as

chlorine-35 and chlorine-37



Isotopes

Isotopes of an element have the same chemicalproperties because they have the same number ofelectrons. When a chemical reaction takes place, itis the electrons that are involved in the reactions.

However isotopes of an element have the slightlydifferent physical properties because they havedifferent numbers of neutrons, hence differentmasses.

The isotopes of an element with fewer neutronswill have: Lower masses • faster rate of diffusion

Lower densities • lower melting and boiling points



Isotopes - Questions

1. Explain what isotopes using hydrogen as anexample.

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

neutrons?

3. State the number of protons, electrons andneutrons in:

a) one atom of carbon-12b) one atom of carbon-14

c) one atom of uranium-235

d) one atom of uranium-238



Isotopes – H/W

Complete Exercise 1, 2, and 3 in thehandbook for next session.

Task: Find out the uses of isotopes inas much detail as possible.

N.B. Please make sure you understand

and write in your own words – DO NOTCOPY out of a text-book.



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 asample of an element

When charged particles

pass through a magnetic

field, the particles are

deflected by the magneticfield, 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 beenplaced in the mass spectrometer, it

undergoes five stages. Vaporisation – the sample has to be ingaseous form. If the sample is a solid orliquid, a heater is used to vaporise some of

the sample.X (s) X (g)

or X (l) X (g)




Ionisation – sample is bombarded by astream of high-energy electrons froman electron gun, which ‘knock’ an

electron from an atom. This produces apositive 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 Spectrometer

If all the ions are travelling at the samevelocity and carry the same charge, the

amount of deflection in a given magnetic fielddepends upon the mass of the ion.

For a given magnetic field, only ions with aparticular relative mass (m ) to charge (z )

ration – the m/z value – are deflectedsufficiently to reach the detector.



Mass Spectrometer

Detection – ions that reach the detectorcause electrons to be released in an ion-current detector

The number of electrons released, hence thecurrent produced is proportional to thenumber of ions striking the detector.

The detector is linked to an amplifier andthen to a recorder: this converts the currentinto a peak which is shown in the massspectrum.



Atomic Structure – Mass Spec

Name the five stages which the sampleundergoes in the mass spectrometer and

make brief notes of what you rememberunder each stage.

Complete Exercise 4, 5 and 6 in thehandbook. Any incomplete work to be

completed and handed in for next session.Card Sort Activity ???



Atomic Structure – Mass Spec

Isotopes of boron

m/z value 10 11

Relativeabundance %

18.7 81.3

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

(18.7 + 81.3)

= 187 + 894.3

100

= 1081.3 = 10.8

100



Mass Spectrometer – Questions

Complete Exercise 7 14



Energy Levels

Electrons go in shells or energy levels.The energy levels are called principleenergy levels, 1 to 4.

The energy levels contain sub-levels.

Principleenergy level

Number ofsub-levels

1 1

2 2

3 3

4 4

These sub-levels

are assigned the

letters, s, p, d, f



Energy Levels

Each type of sub-level can hold adifferent maximum number of electron.

Sub-levelMaximumnumber ofelectrons

s 2

p 6

d 10

f 14



Energy Levels

The energy of the sub-levels increasesfrom 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 theelectrons fill the sub-levels?



Energy Levels

Let’s take a look at the Periodic Table tosee how this fits in.



Electronic Structure

So how do you write it?

1s2

Energy levelSub-level

Number of

electrons

Example

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



Electronic StructureThe electronic structure follows a pattern – the orderof filling the sub-levels is 1s, 2s, 2p, 3s, 3p…

After this there is a break in the pattern, as that the4s fills before 3d.

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

• This is because the 4s

sub-level is oflower energy than the

3d sub-level.




The order in this the energy levels arefilled is called the Aufbau Principle.

Example (Sodium – 2, 8, 1)




There are two exceptions to the Aufbauprinciple.

The electronic structures of chromium andcopper do not follow the pattern – they areanomalous.

Chromium – 1s2, 2s2, 2p6, 3s2, 3p6, 3d5, 4s1

Copper – 1s2, 2s2, 2p6, 3s2. 3p6, 3d10, 4s1

Write the electronic configuration for the following elements:

a) hydrogen c) oxygen e) copper

b) carbon d) aluminium f) fluorine



Electronic Structure – of ions

When an atom loses or gains electronsto 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 2p6

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 3s1 3p6 3d2 4s2

Ti atom Ti2+ ion

Cr atom Cr3+ ion

1s2 2s2 2p6 3s1 3p6 3d2

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

Chromium (Cr) - loss of 3 e-



Electronic Structure - Questions

Give the full electronic structure of thefollowing poisitve ions:

a) Mg2+ b) Ca2+ c) Al3+

Give the full electronic structure of thenegative ions:

a) Cl- b) Br- c) P3-



Electronic Structure - Questions

Copy and complete the following table:

Atomicno.

Massno.

No. ofprotons

No. ofneutrons

No. ofelectrons

Electronicstructure

Mg 12 1s2

2s2

2p6

3s2

Al2+ 27 10

S2- 16 16

Sc3+

21 45

Ni2+ 30 26



Orbitals

The energy sub levels are made up oforbitals, each which can hold a maximum of 2electrons.

Different sub-levels have different number oforbitals:

Sub-levelNo. of

orbitals

Max. no. of

electronss 1 2

p 3 6

d 5 10

f 7 15



Orbitals

The orbitals in different sub-levels havedifferent shapes:

• s orbitals

1s 2s• p orbitals



Orbitals

Within a sub-level, the electrons occupyorbitals as unpaired electrons rather thanpaired electrons. (This is known as Hund’s

Rule).We use boxes to represent orbitals:

1s

2s

2p

Electronic structure ofcarbon, 1s2, 2s2, sp2



OrbitalsThe arrows represent the electrons in theorbitals.

The direction of arrows indiactes the spin of

the electron.Paired electrons will have opposite spin, asthis reduced the mutual repulsion betweenthe paired electrons.

Electronic structure ofcarbon, 1s2, 2s2, 2p2

1s

2s

2p



OrbitalsUsing boxes to represent orbitals, give the fullelectronic structure of the following atoms:

a) lithium b) fluorine c) potassium

d) nitrogen e) oxygen

1s

2s

2p






Electronic structure oflithium: 1s2, 2s1

1s

2s

2p






Electronic structure offluorine: 1s2, 2s2

1s

2s

2p






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

1s

2s

2p

3s

3p

4s






Electronic structure ofnitrogen: 1s2, 2s2, 2p3

1s

2s

2p






Electronic structure ofoxygen: 1s2, 2s2, 2p4

1s

2s

2p



Ionisation EnergyIonisation of an atom involves the loss of anelectron to form a positive ion.

The first ionisation energy is defined asthe energy required to remove one electronfrom a gaseous electron.

The first ionisation energy of an atom can berepresented by the following general

equation:X(g) X+ + e- ΔH +ve

Since all ionisations requires energy, they areendothermic processes and have a positive

enthalpy change ( ΔH) value.



Ionisation Energy

The value of the first ionisation energydepends upon two main factors:

The size of the nuclear charge

The energy of the electron that hasbeen removed (this depends upon its distance from the nucleus)



Ionisation Energy As the size of the nuclear charge increases the forceof the attraction between the negatively chargedelectrons and the positively charged nucleusincreases.

+ +

Smallnuclearcharge

Largenuclearcharge

Small force

of attraction

Smallerionisationenergy

Large force

of attraction

Greaterionisationenergy



Ionisation energy

As the energy of the electron increases, the electronis farther away from the nucleus. As a result theforce of attraction between the nucleus and the

electron decreases.

+Electrons closer topositive nucleus

Large force ofattraction

Greaterionisationenergy

Electrons furtheraway from

positive nucleus

Small force ofattraction

Smallerionisationenergy

+



Ionisation energy - Questions

Write an equation to represent the firstionisation of:

a) aluminiumb) lithium

c) sodium



Trends across a Period

Going 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 nuclearcharge 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 removedfrom Al is in a higher sub-level: the electron removedfrom Al is a 3p electron, whereas that removed fromMg is a 3s.

Electronic structure Ionisation energy/kJ mol-1

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

Mg 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 (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 Pis an unpaired electron.

When the electrons are paired the extra mutualrepulsion results in less energy being required toremove an electron, hence a reduction in the ionisationenergy.

3s

3p

Phosphorus

3s

3p

Sulphur



Trends across a Period - Questions

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

Look at the table below and point out where the break inthe 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/kJmol-1

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

Mg 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 aspossible

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)

F i r s t i o n i s

a t i o n

e n e r g y / k J

m o l - 1

H

He

Li

Be

B

C

N

O

F

Ne

Na

Mg

AlSi

P

S

Cl Ar

K

Ca



Trends down a Group+

+

+

+

D o

wn t h e G r o u

p

Ionisation energy decreases goingdown a Group.

Going down a Group in the PeriodicTable, the electron removed during

the first ionisation is from a higherenergy level and hence it is furtherfrom the nucleus.

The nuclear charge also increases,

but the effect of the increasednuclear charge is reduced by theinner electrons which shield the outerelectrons.



Ionisation energy - Questions

1. Explain why sodium has a higher firstionisation energy than potassium.

2. Explain why the first ionisation energyof boron is less than that of beryllium.

3. Why does helium have the highestfirst ionisation energy of all the

elements?4. Complete Tasks

atomic structure chemistry a level

Documents