1. Electronic Structure of Atoms & Periodic Table
CHEMISTRY - DMCU 1233
Fakulti Kejuruteraan Mekanikal, UTeM
Lecturer:
IMRAN SYAKIR BIN MOHAMAD
MOHD HAIZAL BIN MOHD HUSIN
NONA MERRY MERPATI MITAN
Electronic Structure of Atoms & Periodic Table
Chapter 4

2. 4.1 HISTORY OF ATOMIC MODEL
Contributor
Model
Explanation
John Dalton
(1805)
Billiard Ball Model
Daltons atomic model was represented as a small united ball similar to a very tiny ball.
J. J Thomson
(1897)
Plum Pudding Model
Thomson discovered the electron, a negatively charged particle. The atom was described as a sphere of positive charge with electrons embedded in it.
Ernest Rutherford (1911)
Solar System Model
Rutherford discovered the proton, a positively charged particle in an atom. The proton and most of the mass of the atom were concentrated in the central region called the nucleus. The electrons moved in the spherical space outside the nucleus.
Neils Bohr (1913)
Bohr Model
According to Bohr, the electrons in an atom were not randomly distributed around the atomic nucleus, but moved around the nucleus in fixed orbits (shell). Each orbit formed a circle and had a fixed distance from the nucleus.

3. 4.1 HISTORY OF ATOMIC MODEL
Plum Pudding Model
Billiard Ball Model
Bohr Model
Solar System Model

4. A quantum number describes the energies of electrons in atoms
4.2 Quantum Numbers
A quantum number describes the energies of electrons in atoms
The Bohr model was 1-D model that used one quantum number to describe the distribution of electrons in the atom that representative of the size of the orbit, which was described by the principal quantum number (n).
Meanwhile Schrödinger's model allowed the electron to occupy in 3-D space to describe the orbitals in which electrons can be found.
Each electron in an atom is described by four different quantum numbers.
The first three quantum number from Schrödinger's wave equations are the principal (n), angular (l), and magnetic (ml) quantum numbers describe the size, shape, and orientation in space of the orbitals on an atom.
The fourth quantum number spin (ms) specifies how many electrons can occupy that orbital.

5. 4.2 Quantum Numbers Principal quantum number – ( n )
Angular momentum quantum number – ( l )
Magnetic quantum number – ( ml )
Spin quantum number – ( ms )

6. Quantum Numbers (n, l, ml, ms)
Principal quantum number n
n = 1, 2, 3, 4, ….
Specifies the energy of an electron and the size of the orbital (the distance from the nucleus of the peak in a radial).
All orbitals that have the same value of n are said to be in the same shell (level)
The total number of orbitals for a given n value is n2
n=3
n=2
n =1

7. Quantum Numbers (n, l, ml, ms)
Angular momentum quantum number l
for a given value of n, l = 0, 1, 2, 3, … n-1
Specifies the shape of an orbital with a particular principal quantum number.
The secondary quantum number divides the shells into smaller groups of orbitals called subshells (sublevels).
l = s orbital
l = p orbital
l = d orbital
l = f orbital
n = 1, l = 0
n = 2, l = 0 or 1
n = 3, l = 0, 1, or 2
Shape of the “volume” of space that the e- occupies

8. l = 0 (s orbitals)
l = 1 (p orbitals)

9. l = 2 (d orbitals)

10. l = 3 (f orbitals)

11. Quantum Numbers (n, l, ml, ms)
Magnetic quantum number ml
for a given value of l ml = -l, …., 0, …. +l
Specifies the orientation in space of an orbital of a given energy (n) and shape (l).
This number divides the subshell into individual orbitals which hold the electrons.
There are 2l+1 orbitals in each subshell.
for l = 0 (s orbital) ml = 0
if l = 1 (p orbital), ml = -1, 0, or +1
if l = 2 (d orbital), ml = -2, -1, 0, +1, or +2
orientation of the orbital in space

12. ml = -1
ml = 0
ml = 1
ml = -2
ml = -1
ml = 0
ml = 1
ml = 2

13. ORIENTATION OF THE ORBITAL
IN SPACE

14. Quantum Numbers (n, l, ml, ms)
Spin quantum number ms
ms = +½ or -½
Specifies the orientation of the spin axis of an electron.
An electron can spin in only one of two directions (sometimes called up and down).
ms = +½
ms = -½
Experimental arrangement for demo the spinning motion of electrons
Q & A

15. (orientation of the subshell's shape)
Quantum number
Symbol
Meaning
Range of values
Value examples
principal n
shell
1 ≤ n
n = 1, 2, 3, …
angular momentum ℓ
subshell
(s orbital is listed as 0,
p orbital as 1 )
0 ≤ ℓ ≤ n − 1
for n = 3: ℓ = 0, 1, 2 (s, p, d)
magnetic mℓ
Orbital
(orientation of the subshell's shape)
−ℓ ≤ mℓ ≤ ℓ
for ℓ = 2: mℓ = −2, −1, 0, 1, 2
spin ms
spin of the electron
(−½ = "spin down",
+½ = "spin up")
−s ≤ ms ≤ s
for an electron
s = ½, so ms = −½, +½

16. Relation between quantum number, atomic orbital and
number of an electron n l ml
Number of orbitals
Orbital Name
Number of electrons 1 1s 2 2s
-1, 0, +1 3 2p 6 3s 3p
-2, -1, 0, +1, +2 5 3d 10 4 4s 4p 4d
-3, -2, -1, 0, +1, +2, +3 7 4f 14

17. Quantum Numbers (n, l, ml, ms)
Existence (and energy) of electron in atom is described by its unique Quantum Numbers
Pauli exclusion principle
(Wolfgang Pauli, Nobel Prize 1945)
No two electrons in the same atom can have identical values for all four of their quantum numbers
Two electrons in the same orbital must have opposite spins.
Because an electron spins, it creates a magnetic field, which can be oriented in one of two directions.
For two electrons in the same orbital, the spins must be opposite to each other; the spins are said to be paired

18. Diamagnetic Paramagnetic all electrons paired unpaired electrons 2p 2p
The substances are not attracted to magnets and are said to be diamagnetic.
Atoms with more electrons that spin in one direction than another contain unpaired electrons. These substances are weakly attracted to magnets and are said to be paramagnetic.
Diamagnetic
Paramagnetic
all electrons paired
unpaired electrons 2p 2p

19. Quantum Numbers (n, l, ml, ms)
Shell – electrons with the same value of n
Subshell – electrons with the same values of n and l
Orbital – electrons with the same values of n, l, and ml
How many electrons can an orbital hold?

20. How many 2p orbitals are there in an atom?
How many electrons can be placed in the 3d subshell?
Q & A

21. 4.3 Electron configuration
Electron configuration is how the electrons are distributed among the various atomic orbitals in an atom.
number of electrons in
the orbital or subshell
1s1
principal quantum
number n
angular momentum
quantum number l
Orbital diagram is shows the spin of the electron
1s1 H

22. Order of orbitals (filling) in multi-electron atom
“Fill up” electrons in lowest energy orbitals
Aufbau principle - electrons fill orbitals starting at the lowest available (possible) energy states before filling higher states.
1s < 2s < 2p < 3s < 3p < 4s < 3d < 4p < 5s < 4d < 5p < 6s < 4f < 5d < 6p < 7s < 5f

23. Outermost subshell being filled with electrons
The order in which the electrons are filled in can be read from the periodic table in the following fashion

24. C 6 electrons
C 1s22s22p2
B 5 electrons
B 1s22s22p1
Li 3 electrons
Li 1s22s1
H 1 electron
H 1s1

25. The most stable arrangement of electrons in subshells is the one with the greatest number of parallel spins (Hund’s rule).
Ne 10 electrons
Ne 1s22s22p6
F 9 electrons
F 1s22s22p5
O 8 electrons
O 1s22s22p4
N 7 electrons
N 1s22s22p3

26. What is the electron configuration of Mg?
What are the possible quantum numbers for the last (outermost) electron in Cl?

27. Electron Configurations of Cations and Anions
Of Representative Elements
Na [Ne]3s1
Na+ [Ne]
Atoms lose electrons so that cation has a noble-gas outer electron configuration.
Ca [Ar]4s2
Ca2+ [Ar]
Al [Ne]3s23p1
Al3+ [Ne]
H 1s1
H- 1s2 or [He]
Atoms gain electrons so that anion has a noble-gas outer electron configuration.
F 1s22s22p5
F- 1s22s22p6 or [Ne]
O 1s22s22p4
O2- 1s22s22p6 or [Ne]
N 1s22s22p3
N3- 1s22s22p6 or [Ne]

28. Cations and Anions Of Representative Elements+1 +2 +3 -3 -2 -1

29. Na+, Al3+, F-, O2-, and N3- are all isoelectronic with Ne
Na+: [Ne]
Al3+: [Ne]
F-: 1s22s22p6 or [Ne]
O2-: 1s22s22p6 or [Ne]
N3-: 1s22s22p6 or [Ne]
Na+, Al3+, F-, O2-, and N3- are all isoelectronic with Ne
What neutral atom is isoelectronic with H- ?

30. Electron Configurations of Cations of Transition Metals
When a cation is formed from an atom of a transition metal, electrons are always removed first from the ns orbital and then from the (n – 1)d orbitals.
Fe: [Ar]4s23d6
Mn: [Ar]4s23d5
Fe2+: [Ar]4s03d6 or [Ar]3d6
Mn2+: [Ar]4s03d5 or [Ar]3d5
Fe3+: [Ar]4s03d5 or [Ar]3d5

31. Ion charges

32. Q & A session
Name the orbital described by the following quantum numbers :
n = 3, l = 0
n = 3, l = 1
n = 3, l = 2
n = 5, l = 0

33. Give the n and l values for the following orbital
Q & A session
Give the n and l values for the following orbital
    a. 1s     b. 3s     c. 2p     d. 4d     e. 5f
What and the possible ml values for the following types of orbital?
    a. s     b. p     c. d     d. f

34. How many possible orbital are there for n = a. 4 b. 10
Q & A session
How many possible orbital are there for n =
    a. 4     b. 10
How many electrons can inhabit all of the n = 4 orbital?
Place the following orbital in order of increasing energy:
    1s, 3s, 4s, 6s, 3d, 4f, 3p, 7s, 5d, 5p

35. Write electron configurations for the following atoms: a. H b. Li+
Q & A session
Write electron configurations for the following atoms:
    a. H
    b. Li+
    c. N
    d. F-
    e. Ca

36. Q & A session
Draw an orbital diagrams for atoms with the following electron configurations:
1s22s22p63s23p3

38. When the Elements Were Discovered

39. 4.4 HISTORY OF THE PERIODIC TABLE
Antoine Lavoisier (1743–1794)
Classify elements into four groups including light and heat, into metals and non-metals.

40. 4.4 HISTORY OF THE PERIODIC TABLE
Johann Dobereiner (1780–1849)
The first significant groupings of elements by place certain elements in groups of three known as The Law of Triads.
Founded that strontium had about the average properties of calcium and barium, and grouped these three together.
Several more triad groups, including the halogen triad of chlorine, bromine, and iodine, and the alkali metal triad of lithium, sodium, and potassium.
However, due to the inaccuracy of many measurements, including atomic weight, the relationship between large element groups could not be exacted

41. 4.4 HISTORY OF THE PERIODIC TABLE
John Newlands (1837–1898)
arranged known elements horizontally in the ascending order of their atomic masses
Each row consisted of seven elements.
Founded that elements with similar properties
repeated at every eighth element.
This arrangement was known as the Law of Octaves
However, this law was only obeyed by the first 17 elements.
There were no positions allocated for elements yet to be discovered

42. 4.4 HISTORY OF THE PERIODIC TABLE
Lothar Meyer (1830–1895)
Plotted a graph of atomic volume against atomic mass for all known elements.
Founded that elements with the same chemical properties occupied the same relative positions on the curve.
Showed that the properties of the elements were in a periodic pattern with their atomic masses.
Proved that the properties of the elements recur periodically.

43. 4.4 HISTORY OF THE PERIODIC TABLE
Dmitri Mendeleev (1834–1907)
Showed that the properties of elements changed periodically with their atomic mass.
Arranged the elements in the order of increasing atomic mass and grouped them according to similar chemical properties.
Able to predict the properties of undiscovered elements and left gap for these elements
For examples correctly predicted the properties of the elements gallium, scandium and germanium which were only discovered later
Mendeleevs table was used as a blueprint for the modern periodic table
Mendeleev’s periodic table

44. 4.4 HISTORY OF THE PERIODIC TABLE
Henry J. G. Moseley (1887–1915)
Based on the x-ray spectrum of elements studies, he concluded that the proton numbers should be used as a basis for the periodic change of chemical properties instead of the atomic mass.
Rearranged the elements in the ascending order of their proton numbers
Similar to Mendeleev, Moseley left gaps for elements yet to be discovered.
produced a periodic table which was almost the same as Mendeleevs periodic table.
Due to Moseley’s work, the periodic table was successfully developed and being used today.
The modern periodic table is based on the arrangement of elements in the ascending order of their proton numbers.

45. 4.4 HISTORY OF THE PERIODIC TABLE
Glenn Seaborg
discovered that the transuranium elements that have atomic numbers from 94 to 102, resulting in the redesign of the periodic table
Technically, both the lanthanide and actinide series of elements are to be placed between the alkaline earth metal and the transition metal.
However, by doing this, the periodic table would be too wide.
Thus, the lanthanide and actinide series of elements were placed under the rest of the periodic table.
Dr Seaborg and his colleagues were also responsible for identifying more than 100 isotopes of elements.

46. 4.5 MODERN PERIODIC TABLE
The periodic table is a systematic classification of elements whereby elements with the same chemical properties are placed in the same group.
The elements in the periodic table are arranged in rows called the periods and columns which are known as the groups

47. 4.5 MODERN PERIODIC TABLE Groups
There are 18 groups of elements in the periodic table.
Some of these groups have special names:
(a) Group 1 elements are called alkali metals.
(b) Group 2 elements are called alkaline earth metals.
(c) Group 3 to Group 12 elements are known as transition elements.
(d) Group 17 elements are called halogens.
(e) Group 18 elements are called noble gases.
Each member of a group shows similar chemical properties although their physical properties such as density, melting point and colour show a gradual change when descending the group.

48. 4.5 MODERN PERIODIC TABLE Periods
There are seven rows from period 1 to period 7.
The elements are arranged horizontally in the ascending order of their proton numbers in the periodic table.
The position of the period of an element in the periodic table is determined by the number of shells occupied with electrons in the atom.
Period 1 has 2 elements only H and He,
Periods 2 and 3 have 8 elements each.
Periods 4 and 5, they have 18 elements each and they are called the long periods.
Period 6 has 32 elements whereas the elements with proton number 58 to 71 are separated and are grouped below the periodic table known as the Lanthanide Series.
Period 7 has 32 elements, the elements with proton number 90 to 103 are grouped below the periodic table known as the Actinide Series.

49. Periodic Classification
Classification as metals and non-metals
(a) Metals – a good conductor of heat and electricity.
(b) Non-metals - a poor conductor of heat and electricity.
(c) Metalloids – a intermediate between metal and non-metal properties

50. Periodic Classification
Ground State Electron Configurations of the Elements
The similarity of the outer electron configuration (same type of valence electrons) makes the elements in the same group resemble one another in chemical behavior.
ns2np6
ns2np1
ns2np2
ns2np3
ns2np4
ns2np5
ns1
ns2 d1 d5
d10 4f 5f

51. Periodic Classification
Classification based on subshell filled with electron

52. Periodic Classification
Representative elements (incompletely filled s and p subshells)
Transition metals (incompletely filled d subshells)
Noble gases (completely filled p subshells)
Actinides (incompletely filled 5f subshells)
Lanthanides (incompletely filled 4f subshells)
Zn, Cd, Hg (neither representative element nor transition metals)

53. the periodic table of elements
Periodic Trends
Periodic trends are the tendencies of certain elemental characteristics to increase or decrease as one progresses along a row or column of
the periodic table of elements
Elemental characteristics
Atomic radius
Ionization energy
Electron affinity
Electronegativity
Metallic properties
Non-metallic properties

54. Periodic Trends Atomic radius
The atomic radius is the distance between an atom's nucleus and its valence electrons in an atom.
The atomic radius tends to decrease as one progresses across a period from left to right because the effective nuclear charge increases, thereby attracting the orbiting electrons and reducing the radius.
The atomic radius usually increases while going down a group due to the addition of a new energy level (shell).

55. Periodic Trends Ionic Radius
The ionic radius is different from the atomic radius of an element.
Positive ions are smaller than their uncharged atoms.
Negative ions are larger than their atoms.

56. Periodic Trends Ionization energy
The ionization potential is the minimum amount of energy required to remove one electron from each atom.
Ionization energies increase moving from left to right across a period (decreasing atomic radius).
Ionization energy decreases moving down a group (increasing atomic radius).

57. Periodic Trends Electron affinity
The electron affinity described as the energy gained by an atom when an electron is added.
Electron affinities becoming increasingly from left to right
Electron affinities change little moving down a group, however they do generally become slightly more positive (less attractive toward electrons)

58. Periodic Trends Electronegativity
Electronegativity refers to the ability of an atom to attract the electrons of another atom to it when those two atoms are associated through a bond.
Electronegativity generally increases moving across a period and decreases moving down a group.
Electronegativity plays a very large role in the processes of Chemical Bonding.

59. Periodic Trends Metallic properties
Metallic property decreases across a period with increase in number of valence electrons as well as a decrease in atomic radius, and it increases down the group with increase in number of shells and atomic radius.
Non-metallic properties
Non-metallic property increases across a period and decreases down the group due to the same reason