1. Topic 2 Atomic structure
IB CHEMISTRY
Topic 2 Atomic structure
Higher level

2. 2.1 The nuclear atom
OBJECTIVES
• Atoms contain a positively charged dense nucleus composed of protons and neutrons (nucleons).
• Negatively charged electrons occupy the space outside the nucleus.
• The mass spectrometer is used to determine the relative atomic mass of an element from its isotopic composition.
• Use of the nuclear symbol notation to deduce the number of protons, neutrons and electrons in atoms and ions.
• Calculations involving non-integer relative atomic masses and abundance of isotopes from given data, including mass spectra.

3. Atomic Structure Atoms are very small ~ 10-10 metres
All atoms are made up of three sub-atomic particles: protons, neutrons, and electrons
Label this diagram:
The protons and neutrons form a small positively charged nucleus
The electrons are in energy levels outside the nucleus

4. Subatomic particles
The actual values of the masses and charges of the sub-atomic particles are shown in your data booklet:
A meaningful way to consider the masses of the sub-atomic particles is to use relative masses

5. Element (X)
Atomic number (Z) is the number of protons in the nucleus of an atom. It is also known as the proton number. 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. Some periodic tables have Z above A. Remember A will always be the biggest number.

6. Problem:
Calculate the number of protons and neutrons in:
Z = number of protons = 17 protons
A = number of protons and neutrons = 35
Number of neutrons = A - Z= = 18 neutrons

7. 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 protons
17 electrons
18 neutrons
17 protons
17 electrons
20 neutrons
The isotopes of chlorine are often referred to as chlorine-35 and chlorine-37

8. Some Isotopes of Carbon

9. Properties of 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
Lower densities
Lower melting and boiling points

10. Radioisotopes
Radioisotopes are isotopes that have unstable nuclei and therefore emit radiation when then break up.
Radioisotopes are very useful in society:
14C is used in radiocarbon dating14C is used in radiocarbon dating
Detecting gas leaks
Industrial quality control
60Co is used in radiotherapy
131I and 125I are used a medical tracers
Nuclear power
Radioisotopes can also be very dangerous to living things:
Radioactive contamination of the environment
Radiation poisoning

11. Industrial use: detecting blockages in underground pipes
A radioactive isotope which is a gas gets passed down the pipe, where it concentrates the blockage is present.

12. Industrial use: Quality Control
The radioactive isotope is used as a source of radiation and the amount penetrating the material gives a measure of it’s thickness.

13. Radiotherapy
A cobalt-60 source can be rotated around the patient. The gamma rays emitted are focussed on the tumour. Healthy surrounding tissue receives a much smaller dose. The cells in the tumour are damaged while surrounding tissue is not.

14. A radioactive sample can be swallowed
A radioactive sample can be swallowed. The chemical chosen will be one that concentrates in a particular area. For example, cancer of the thyroid can be treated using iodine-131.

15. Mass Spectrometer
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.

16. Ar of boron = (amu1 x %1) + (amu2 x %2) total %
Problem1: Determine the relative atomic mass of boron from the following spectrograph:
m/z value 11 10
Relative abundance %
18.7
81.3
Ar of boron = (amu1 x %1) + (amu2 x %2)
total %
= (11 x 18.7) + (10 x 81.3)
( ) =
100
Ar = = 10.2

17. Problem 2: A mass spec chart for a sample of neon shows that it contains 90.9% 20Ne, 0.17% 21Ne, and 8.93% 22Ne. Calculate the relative atomic mass of neon.
Ar of neon = (amu1 x %1) + (amu2 x %2) + (amu3 x %3) total %
= (90.9 x 20) + (0.17 x 21) + (8.93 x 22) =
100
Ar = =

18. 2.2 Electron configuration
OBJECTIVES
• Emission spectra are produced when photons are emitted from atoms as excited electrons return to a lower energy level. • The line emission spectrum of hydrogen provides evidence for the existence of electrons in discrete energy levels, which converge at higher energies. • The main energy level or shell is given an integer number, n, and can hold a maximum number of electrons, 2n2. • A more detailed model of the atom describes the division of the main energy level into s, p, d and f sub-levels of successively higher energies. • Sub-levels contain a fixed number of orbitals, regions of space where there is a high probability of finding an electron. • Each orbital has a defined energy state for a given electronic configuration and chemical environment and can hold two electrons of opposite spin. • Description of the relationship between colour, wavelength, frequency and energy across the electromagnetic spectrum. • Distinction between a continuous spectrum and a line spectrum. • Description of the emission spectrum of the hydrogen atom, including the relationships between the lines and energy transitions to the first, second and third energy levels. • Recognition of the shape of an s atomic orbital and the px, py and pz atomic orbitals. • Application of the Aufbau principle, Hund’s rule and the Pauli exclusion principle to write electron configurations for atoms and ions up to Z = 36.

19. Electromagnetic Spectrum
short
long

20. Wavelength (λ) In the data booklet: E=hν and c=νλ
(where ν is the frequency, and h and c are constants)
(ν is the small greek letter N, not v for velocity, pronounced nu, λ
is pronounced lambda)
It follows then that the shorter the wavelength, the higher the frequency of the wave, and the more energy it contains.

21. Emission spectra
When electrons are excited to a higher energy level, and then return to a lower energy level, they release a photon of a specific energy, as shown by a specific frequency of light.

22. Spectroscopy Continuous spectrum Emission (line) spectrum
Absorption spectrum

23. Spectral line series

24. Emission Spectrum of Hydrogen - convergence
Electrons moving back to the lowest energy states and over the longest distances release the highest E (short λ). In each series the lines converge meaning higher levels/shells get closer together.

25. Actual values

26. Convergence
Outer shells become closer together, so spectral lines get closer together – called convergence.

27. Spectral fingerprint of the elements

28. Spectrophotometer with discharge lamps

29. Periodic table of element emission and absorption spectra

30. Electron Shells
Although simplistic, a useful way to look at shells is to use the periods in the Periodic table.

31. Electron Configuration
Electrons go in shells or energy levels. The energy levels are called principle energy levels, 1 to 4. The maximum number of electrons an energy level (n) can hold is 2n2.
The energy levels contain sub-levels.
Principle energy
level
Maximum number of electrons
Number of sub-levels 1 2 8 3 18 4 32
These sub-levels are assigned the letters,
s, p, d, f

32. Maximum number of electrons
Sublevels
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 f 14

33. Electron Configuration
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.

34. Electron Configuration
Let’s take a look at the Periodic Table to see how this fits in.

35. (One of these needs to be memorized)

36. Electron Configuration
So how do you write it?
1s2
Example
For magnesium:
1s2, 2s2, 2p6, 3s2
Energy level
Number of
electrons
Sub-level

37. Electron Configuration
The electronic configuration 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 the 4s fills before 3d (The electrons fill up the lower energy sub-levels first)
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.

38. Electrons and Sub-Levels

39. Writing Electronic Configurations
The order in which the energy levels are filled is called the Aufbau Principle.
Example (Sodium: 2, 8, 1)

40. Writing Electronic Configurations
There are two exceptions to the Aufbau principle.
The electronic configurations of chromium and copper do not follow the pattern – they are anomalies!
Chromium – 1s2, 2s2, 2p6, 3s2, 3p6, 3d5, 4s1
Copper – 1s2, 2s2, 2p6, 3s2. 3p6, 3d10, 4s1

41. Writing Electronic Configurations for ions
When an atom loses or gains electrons to form an ion, the electronic configuration changes:
Positive ions: formed by the loss of e-
Negative ions: formed by the gain of e-
1s2 2s2 2p6 3s1  1s2 2s2 2p6
Na atom
Na+ ion
1s2 2s2 2p4  1s2 2s2 2p5
O atom
O- ion

42. Writing Electronic Configurations for 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-
Chromium (Cr) - loss of 3 e-
1s2 2s2 2p6 3s2 3p6 3d2 4s2  1s2 2s2 2p6 3s2 3p6 3d2
Ti atom
Ti2+ ion
1s2 2s2 2p6 3s2 3p6 3d5 4s1  1s2 2s2 2p6 3s2 3p6 3d3
Cr atom
Cr3+ ion

43. Writing Electronic Configurations – Condensed form
Abbreviations can also be used in electron configuration for simplicity sake.
Titanium (Ti):
1s2 2s2 2p6 3s2 3p6 3d2 4s2  Ar] 3d2 4s2 or [Ar] 4s2 3d2
[Ar] always represents 1s2 2s2 2p6 3s2 3p6
Other noble gases in the VIII group of the periodic table are used as well, such as, [He], [Ar], [Kr], [Xe], etc.

44. 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 d 5 10 f 7 14

45. Energy Levels and Sub-levels
Main energy level
Sub-levels
Max. no. of electron pairs in sub-level
Max. no. of electrons in sub-level
Max. no. of electrons in main level 1 s 2 8 p 3 6 18 d 5 10 4 32 f 7 14

46. Shapes of the orbitals

47. Orbital diagrams
Within a sub-level, the electrons occupy orbitals as unpaired electrons rather than paired electrons and these all spin in the same direction. (This is known as Hund’s Rule).
We use boxes to represent orbitals: 1s 2s 2p    
Electronic configuration of carbon, 1s2, 2s2, 2p2  

48. Orbital diagrams The arrows represent the electrons in the orbitals.
The direction of arrows indicates the spin of the electron.
Paired electrons will have opposite spin, as this reduces the mutual repulsion between the paired electrons (This is known as the Pauli exclusion principle)
Notice how the arrows in each box of the 1s2 & 2s2 are opposite which means opposite spin 1s 2s 2p 
Electron configuration of carbon:
1s2, 2s2, 2p2

49. Problem: Using boxes to represent orbitals, give the full electronic configuration of the following atoms:
a) lithium
b) fluorine
c) potassium
d) nitrogen
e) oxygen 1s 2s 2p

50. Lithium: 1s2, 2s1 1s 2s 2p   
Fluorine: 1s2, 2s2, 2p5 1s 2s 2p 

51. 1s 2s 2p 3s 3p   4s Nitrogen: 1s2, 2s2, 2p3
Potassium: 1s2, 2s2, 2p6, 3s2, 3p6, 4s1
Nitrogen: 1s2, 2s2, 2p3 1s 2s 2p 

52. Oxygen: 1s2, 2s2, 2p4 1s 2s 2p  

53. Orbital diagram review
Electrons enter the lowest energy orbital available (Aufbau principle).
Electrons prefer to occupy orbitals on their own, spin in the same direction, and only pair up when no empty orbitals of the same energy are available (Hund's Rule).
Paired electrons have the opposite spin (Pauli exclusion principle).
In ions, the electrons in the highest energy levels are lost first, but when losing electrons, electrons are lost from 4s before 3d (the energy levels are very close, and when electrons fill them, 4s goes above 3d).

54. 12.1 Electrons in atoms
Higher level
OBJECTIVES
• In an emission spectrum, the limit of convergence at higher frequency corresponds to the first ionization energy. • Trends in first ionization energy across periods account for the existence of main energy levels and sub-levels in atoms. • Successive ionization energy data for an element give information that shows relations to electron configurations. • Solving problems using E=hv. • Calculation of the value of the first ionization energy from spectral data which gives the wavelength or frequency of the convergence limit. • Deduction of the group of an element from its successive ionization energy data. • Explanation of the trends and discontinuities in first ionization energy across a period.

55. Higher level
Ionisation Energy
Ionisation of an atom involves the loss of an electron to form a positive ion.
The first ionisation energy is defined as the energy required to remove one mole of electrons from one mole of atoms of a gaseous element.
The first ionisation energy of an atom can be represented by the following general equation:
X(g)  X+ + e- ΔH: +ve
Since all ionisations requires energy, they are endothermic processes and have a positive enthalpy change (ΔH) value.

56. Emission Spectrum and Ionisation Energy
Higher level
Emission Spectrum and Ionisation Energy
As the electron moves to higher energy states, the spectral lines converge. The upper limit of this convergence is the amount of energy it takes to remove the electron – the ionization energy (IE).

57. Calculations of IE given ν or λ
Higher level
Calculations of IE given ν or λ
Data booklet:
c = νλ and E = hν
c = speed of light = 3.00x108ms-1
h = Planck’s constant = 6.63x10-34Js
ν = frequency (Hz or s-1) pronounced nu
λ = wavelength (m) pronounced lambda
E = energy (J)

58. Higher level
Problem 1: Determine the energy of a photon of red light given the wavelength is 650nm.
c = νλ
3.00x108ms-1 = ν x 650 x 10-9m
∴ ν = 4.62x1014/s
E = hν = 6.63x10-34Js x 4.62x1014/s
= 3.06x10-19J

59. IE = photons/mol x E/photon = 6.02 x 1023 x 2.18x10-18J
Higher level
Problem 2: Determine the first IE of hydrogen given that the shortest-wavelength line in the Lyman series is 91.16nm.
The shortest wavelength will be line at the convergence and hence represent the IE. Units are kJ/mol.
c = νλ
3.00x108ms-1 = ν x x 10-9m
∴ ν = 3.29x1015/s
E = hν = 6.63x10-34Js x 3.29x1015/s
= 2.18x10-18J
IE = photons/mol x E/photon = 6.02 x 1023 x 2.18x10-18J
= J/mol
= 1312 kJ/mol

60. Factors in Ionisation Energy
Higher level
Factors in Ionisation Energy
The value of the first ionisation energy depends upon two main factors:
The size of the nuclear charge
The energy of the electron that has been removed (this depends upon its distance from the nucleus)

61. Higher level
As 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 
Large force of attraction 
Greater ionisation energy
Smaller ionisation energy

62. Higher level
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 further away from positive nucleus  +
Electrons closer to positive nucleus 
Large force of attraction 
Small force of attraction 
Greater ionisation energy
Smaller ionisation energy

63. IE Trends across a Period
Higher level
IE 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 nuclear charge increases. + + + +
Going across a Period

64. IE Trends across a Period
Higher level
IE Trends across a Period
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 configuration
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

65. IE Trends across a Period
Higher level
IE Trends across a Period
The first ionisation energy of Sulfur is less than that of Phosphorus, 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 Sulfur is a paired electron, whereas the 3p electron removed from Phosphorus 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. 3s 3p  
Phosphorus 3s 3p  
Sulphur

66. Now take a look at the graph below:
Higher level
Now take a look at the graph below:
Explain what the graph shows in as much detail as possible
There is one other break in the general pattern going across a Period. What is it and explain why that is.

67. IE Trends down a Group Ionisation energy decreases going down a Group.
Higher level
IE Trends down a Group +
Ionisation energy decreases going down a Group.
Going down a Group in the Periodic Table, the electron removed during the first ionisation 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. + +
Down the Group +

68. Successive Ionisation energy
Higher level
Successive Ionisation energy
Definition: 2nd i.e.
The energy per mole for the process
X+(g) X2+(g) +e-
And so on for further successive ionisation energies

69. Successive Ionisation energy
Higher level
Successive Ionisation 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

70. Successive Ionisation energy
Higher level
Successive Ionisation energy
Large increase between 4th and 3rd shells – electron closer to nucleus
2nd i.e. higher than first – electron has greater pull from nucleus