The document provides a comprehensive overview of atomic structure, detailing components such as the nucleus, isotopes, atomic mass, electron configuration, and quantum theory. It explains key concepts including ionization energy, the arrangement of electrons in energy levels, and the significance of mass spectroscopy in determining atomic and molecular masses. Additionally, it discusses historical figures in the development of atomic theory and the nature of light as it relates to atomic energy levels.

1. Atomic structure
by
Dr. Rania S. Seoudi
1

2. I- the nuclear atom
1- Atom
- Atom: negatively charged electrons orbit around centre positively nucleus.
- Nucleus: positively charged protons and neutrons (except for hydrogen atom no
neutrons).
- Nucleons: protons and neutrons.
- Protons and neutrons are made of small particles called Quarks
- Atom diameter= 1x10-10m
- Nucleus diameter= 1x10-14 to -15m
(nucleus 10000 to 100 000 times smaller than an atom)
2

3. 3
Particles Relative mass ratio Relative charge ratio Actual mass
Proton 1 +1 1.67x10-27kg
Neutron 1 0
electron 5x10-4 -1

4. - Atomic number (Z): is the number of protons in the nucleus.
Number of protons in an atom = number of electrons
- Mass number (A): is the number of protons and neutrons in the nucleus.
Number of neutrons= mass number – atomic number
4

5. - Atom is neutral
- Ions: atom that lost or gained electrons (not neutral)
- Positive ion (cation): atom lost electrons
- Negative ion (anion): atom gained electrons
5

6. 2- Isotopes
- Isotopes: are different atoms of the same element with different mass numbers:
i.e. different numbers of neutrons in the nucleus.
- Isotopes have the same chemical properties (they react in exactly the same
way) but different physical properties (e.g. different melting points and boiling
points due to different atomic masses).
6

7. Discovery of neutrons
- Geiger and Marsden experiment designed by Rutherford.
7

8. Relative atomic masses
- The relative atomic mass (Ar): of an element is the average of the masses of the
isotopes in a naturally occurring sample of the element relative to the mass of 1/12
of an atom of carbon-12.
- How to calculate relative atomic mass:
If natural abundance of Lithium isotops are 6Li=7% and 7Li= 93%.
The relative atomic mass (Ar)=
𝑛𝑎𝑡𝑢𝑟𝑎𝑙 𝑎𝑏𝑢𝑛𝑑𝑎𝑛𝑐𝑒 % 𝑜𝑓6Li+𝑛𝑎𝑡𝑢𝑟𝑎𝑙 𝑎𝑏𝑢𝑛𝑑𝑎𝑛𝑐𝑒 % 𝑜𝑓 7Li
100
=
7𝑥6 +(93𝑥7)
100
= 6.93
8

9. Calculation of percentage composition of
element
- Example: if we are given the relative atomic mass of Lithium =6.93 for the two
naturally occurring isotops 6Li and 7Li.
The present composition will be:
The relative atomic mass (Ar)=
𝑛𝑎𝑡𝑢𝑟𝑎𝑙 𝑎𝑏𝑢𝑛𝑑𝑎𝑛𝑐𝑒 % 𝑜𝑓6Li+𝑛𝑎𝑡𝑢𝑟𝑎𝑙 𝑎𝑏𝑢𝑛𝑑𝑎𝑛𝑐𝑒 % 𝑜𝑓 7Li
100
6.93=
6𝑥+7(100−𝑥)
100
X= 7% = abundance of 6Li
Relative abundance of 7Li= 100-7= 93%
9

10. The mass spectrum of an element and
relative atomic mass (mass spectrometer)
- Mass spectroscopy is the method used for
determining the atomic and molecular masses
- The proportion of each isotope present as one
peak.
- Area under the peak represents the atoms of
each isotops.
10

11. Representation of data in mass spectrometer
- Data represented as peaks.
- The relative atomic mass for the opposite
Figure can be calculated using:
Ar=
78.6𝑥24 +(10.1𝑥25)(11.3𝑥26)
100
= 24.3
11

12. II- Electron configuration
1- the arrangement of electrons in atoms
- Electrons are arranged around nucleus in energy levels (shells) Principle
quantum number.
- Lowest energy level K, then L, M, N, O, P and Q.
- Each level are filled with 2n2 electrons (n is the number of energy level) .
12
Main energy level
number
1 2 3 4 5
Name in letters K L M N O
Maximum number of
electrons
2 8 18 32 50

13. - Rule in filling energy level:
Lowest energy levels are filled first.
- first and second has to be filled first
before moving to the next level.
- Third level filled first with 8e’s then
electrons are put in fourth level.
13

14. Light is a form of energy
- Visible light is part of electromagnetic spectrum
- Electromagnetic radiation is regarded as a waves, particles of this radiation are
called photons .
- Visible light is made of colors of spectrum .
- Arrangement of colors according to increasing energy:
Red<orange< yellow<green< blue< indigo< violet
- The speed of light in vacuum (3.0x108ms-1)
14

15. The electromagnetic spectrum
- When energy is given to electron it is excited then when it loses this energy it
moves back the energy lost as electromagnetic radiation.
- The released radiation is a characteristic for each element.
Frequencyα
1
𝑤𝑎𝑣𝑒𝑙𝑒𝑛𝑔𝑡ℎ
Frequency α energy
Energy = h v =
ℎ
λ
15

16. Evidence for energy levels in atoms
- Light emitted by gas at low pressure and subjected to high voltage is called
emission spectrum.
- emission spectrum consists of sharp, bright lines called line spectrum.
- Emission spectrum consists of all colors merging into each other called
continuous spectrum.
16

17. 17
How an emission spectrum is formed

18. Difference between line spectrum and
continuous spectrum
18

19. Different series of lines
- spectrum fall within Level 2 can be seen by visible eye (Balmer series)
- Transitions down to level 1 occur in ultraviolet region
19

20. Convergence
- Convergence: lines in emission spectrum get close together at higher
frequency /energy.
20

21. 2- Full electron configuration
Sub-energy levels and orbitals
- Main energy levels are made of sub-energy levels (subshells).
- Subshells energy order: s<p<d<f (atomic orbitals)
21

22. Orbitals
- Electrons occupy atomic orbitals around nucleus not orbit like Boher
postulated.
- Orbital: Region of space around nucleus where there is a high probability of
finding electron (it represents a discrete energy level)
22

23. Shapes of orbitals
s orbitals:
Boundary surface diagram: encloses about 90% of the total electron density in an
orbital

24. p orbital: boundary diagram of p like a two lobes on opposite sides of nucleus
It start when n ≥2
For The second main energy level (maximum number of electrons 8) n=2, L=0,
and 1 and ml for L0=0 so one orbital 2s , and for L=1 so m1=-1, 0, 1 three
orbitals there for 2p orbitals are 2px, 2py, and 2pz
px, py, and pz are identical in size, shape, and energy. They differ only in their
orientation with respect to each other.

25. d orbitals: different orientation of d due to different ml values
Lowest value for n for d orbital = 3 so l= and ml= -2, -1, 0, 1, 2 so five d orbitals
(3dxy, 3dyz, 3dx
2-y
2 and 3dz
2)

26. - The third shell (maximum 18 electrons) consists of the 3s, 3p and 3d sub-
levels. The 3s sub-level is just the 3s orbital; the 3p sub-level consists of
three 3p orbitals; and the 3d sub-level is made up of five 3d orbitals.
- The fourth shell (maximum 32 electrons) consists of one 4s, three 4p, five
4d and seven 4f orbitals.
26
Within any subshell, all the orbitals
have the same energy (they are
degenerate) – e.g. the three 2p
orbitals are degenerate and the five
3d orbitals are degenerate.

27. Electron configuration
- It is a code for with the four quantum numbers that describes each element
(fingerprint)

28. Aufbau (building-up) principle
- Aufbau principle: electrons fill sub-levels in the order of increasing energy.
28

29. Abbreviation for writing electronic
configuration
- Example: electronic configuration of germanium (Gr):
1s22s22p63s23p64s23d104p2 (Is abbreviated to [Ar]4s23d104p2)
- All atoms in the same group (vertical column) have the same outer shell
electronic configuration.
- Example: group 16 have outer shell electronic configuration: ns2np4 where n is
the period number
29

30. 30

31. General rules for assigning electrons in atomic orbitals
1- each shell or principle level has n contains n subshells, if n=2 so we have two
subshells
2- each subshell of quantum number l contains 2L+1 orbitals, if l=1 so there is 3p
orbitals
3- no more than two electrons can be placed in each orbitals
4- a quick way to determine max number of electrons formula 2n2
5- principle energy level get closer together as they get further from the nucleus

33. Pauli exclusion principle
- The Pauli exclusion principle: the maximum number of electrons in an orbital
is two. If there are two electrons in an orbital, they must have opposite spin.
- No two electrons in an atom can have the same four quantum numbers
- So for He atom quantum number (1, 0, 0, +1/2) and (1, 0, 0, -1/2)

34. Hund’s rule
- Hund’s rule: electrons fill orbitals of the same energy (degenerate orbitals) so
as to give the maximum number of electrons with the same spin.
- No electron paring occurs till all subshell is filled (parallel spin provide atom
with greater stability)

35. Electronic configuration distribution
Exceptions:
24Cr: [Ar]3d54s1
29Cu: [Ar]3d104s1
35

36. III- electrons in atoms (HL)
1- Ionisation energy and the convergence limit
- The ionisation energy: is the minimum amount of
energy required to remove an electron from a
gaseous atom.
- Ionisation energy can be worked out from known
the frequency of the light emitted at the convergence
limit
- Ionisation energy is energy required for process:
36

37. The relationship between the energy of a
photon and the frequency of electromagnetic
radiation
- The energy (E) of a photon is related to the frequency of the electromagnetic
radiation: E=hv
Where: v is the frequency of the light (Hz or s−1)
h is Planck’s constant (6.63 × 10−34 J s)
- The equation can be used to work out the difference in energy between the
various levels
37

38. The wavelength of the light can be worked out from the frequency using the
equation:
c = vλ
where
λ is the wavelength of the light (m)
c is the speed of light (3.0 × 108 ms–1)
The two equations E = hv and c = vλ can be combined:
E=
ℎ𝑐
λ
38

39. Ionisation energy and evidence for energy
levels and sub-levels
- The first ionisation energy for an element is the energy for the process:
- Definition: the energy required to remove one electron from each atom in one
mole of gaseous atoms under standard conditions
39

40. - The second ionisation energy is:
- The nth ionisation energy is:
40

41. The second ionisation energy is always higher than
the first, and this can be explained in two ways:
- Once an electron has been removed from an atom, a positive
ion is created. A positive ion attracts a negatively charged
electron more strongly than a neutral atom does. More energy is
therefore required to remove the second electron from a
positive ion.
- Once an electron has been removed from an atom, there is less
repulsion between the remaining electrons. They are therefore
pulled in closer to the nucleus. If they are closer to the nucleus,
they are more strongly attracted and more difficult to remove.
41

42. Successive ionisation energies of potassium
- The simple electron arrangement of potassium is 2,8,8,1
- Outermost electron in potassium is furthest from the nucleus and therefore least
strongly attracted by the nucleus and shielded from nucleus attraction effect
- The ionisation energy depends on which main energy level the electron is
removed from.
42

43. Effective nuclear charge
Zeff = Z – σ
where σ (sigma) is called the shielding constant (no. of electrons)
- The effective attraction force felt by the outer electron.
- For potassium: effective nuclear charge = (19+in nucleus -18shielding
electons)= +1
43

44. Explaining ionisation energy graph for silicon
- There is a large jump in the ionisation energy between the fourth and the fifth
ionisation energies, which suggests that these electrons are removed from
different main energy levels.
- Silicon has four electrons in its outer main energy level (shell)
S= 1s22s22p63s23p2.
Fourth ionization= 1s22s22p63s03p0
Fifth = 1s22s22p53s03p0
44

45. Variation in ionisation energy across a period
- Ionization energy is always positive
- Ionization energy:
• increases across the periodic table from left to right.
• decreases moving down the periodic table.
- M(g) → M+
(g) + e- ∆E1= first ionization energy
- First ionization energy usually smaller than
second or third ionization energy

46. Exceptions
- There are two exceptions to the general increase in ionisation energy across a
period.
- First exception: Be= 1s22s2 B= 1s22s22p1
Second exception: N= 1s22s22p3 O= 1s22s22p4
46

47. The transition metals
Exception for “Last in first out” rule
Fe atom= 1s22s22p63s23p64s23d6.
Fe1+= 1s22s22p63s23p64s13d6.
Fe2+= 1s22s22p63s23p63d6.
Fe3+= 1s22s22p63s23p63d5.
47

48. 48

49. 49

50. Further readings
- Democritus and his teacher Leucippus, fifth-century BC Greek philosophers,
first suggesting the idea of the atom as the smallest indivisible particle of which
all matter is made.
- John Dalton (1766–1844) is generally regarded as the founder of modern
atomic theory.
- The electron was discovered in 1897 by J. J. Thompson at the University of
Cambridge, UK.
50

51. Hydrogen isotopes
Hydrogen 1: 1
1
𝐻 Protium
Hydrogen 2: 1
2
𝐻deuterium (D)
Hydrogen-3: 1
3
𝐻tritium (T)
51

52. Nature of science
- Bunsen burner flame (with high temperature) is used in spectroscopic analysis
of substances (yellow color for NaCl salt)
- Neil Bohr: Proposed a model for the atom and used model to explain line
spectra of hydrogen.
52

53. - when the difference in frequency between successive lines is zero, plotting
difference in frequency of successive lines against their frequency gives the
convergence limit.
53