atomic structure chemistry a level
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8/12/2019 Atomic Structure CHEMISTRY A LEVEL
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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?
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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
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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
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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
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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.
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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 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.
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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)
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Mass Spectrometer – 5 Stages
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.
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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• the strength of the magnetic field is
greater
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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.
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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.
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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 ???
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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
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Mass Spectrometer – Questions
Complete Exercise 7 14
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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
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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
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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?
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Energy Levels
Let’s take a look at the Periodic Table tosee how this fits in.
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Electronic Structure
So how do you write it?
1s2
Energy levelSub-level
Number of
electrons
Example
For magnesium:1s2, 2s2, 2p6, 3s2
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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.
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Electronic Structure
The order in this the energy levels arefilled is called the Aufbau Principle.
Example (Sodium – 2, 8, 1)
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Electronic Structure
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
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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-
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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 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-
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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-
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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
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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
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Orbitals
The orbitals in different sub-levels havedifferent shapes:
• s orbitals
1s 2s• p orbitals
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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
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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
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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
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OrbitalsUsing boxes to represent orbitals, give the fullelectronic structure of the following atoms:
a) lithium b) fluorine c) potassium
d) nitrogen e) oxygen
Electronic structure oflithium: 1s2, 2s1
1s
2s
2p
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OrbitalsUsing boxes to represent orbitals, give the fullelectronic structure of the following atoms:
a) lithium b) fluorine c) potassium
d) nitrogen e) oxygen
Electronic structure offluorine: 1s2, 2s2
1s
2s
2p
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OrbitalsUsing boxes to represent orbitals, give the fullelectronic structure of the following atoms:
a) lithium b) fluorine c) potassium
d) nitrogen e) oxygen
Electronic structure of potassium:1s2, 2s2, 2p6, 3s2, 3p6, 4s1
1s
2s
2p
3s
3p
4s
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OrbitalsUsing boxes to represent orbitals, give the fullelectronic structure of the following atoms:
a) lithium b) fluorine c) potassium
d) nitrogen e) oxygen
Electronic structure ofnitrogen: 1s2, 2s2, 2p3
1s
2s
2p
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OrbitalsUsing boxes to represent orbitals, give the fullelectronic structure of the following atoms:
a) lithium b) fluorine c) potassium
d) nitrogen e) oxygen
Electronic structure ofoxygen: 1s2, 2s2, 2p4
1s
2s
2p
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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.
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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)
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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
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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
+
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Ionisation energy - Questions
Write an equation to represent the firstionisation of:
a) aluminiumb) lithium
c) sodium
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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
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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 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
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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 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
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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
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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 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
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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.
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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