Video Lesson

Lesson Objective

At the end of the lesson, you will be able to :

• Correlate the electron configurations of elements with their positions inthe periodic table.
• Give a reasonable explanation for the shape of the periodic table.
• Classify elements as representative, transition and inner-transition elements.
• Explain the general trends in atomic radius, ionization energy, electron affinity, electronegativity, and metallic character of elements within a period and group of the periodic table.
• Write the advantages of the periodic classification of elements.
• Demonstrate periodic law and how electronic configurations of atoms are related to the orbital diagrams
• Explain periodic trend describe scientific enquiry skills along this unit: inferring, predicting, classifying, comparing and contrasting, communicating, asking questions and making generalizations.

Brainstorming Question

• How do the Aufbau principle, Hund’s rule, and the Pauli Exclusion Principle collectively determine the electronic configuration of an atom?
• What is the significance of the Pauli Exclusion Principle in determining the electronic configuration of elements?
• How does the electronic configuration of an atom influence its chemical reactivity and bonding behavior?
• Compare the electronic configurations of ions to their neutral atoms. How does this affect their chemical properties?
• How do atomic radius, ionization energy, and electron affinity vary across periods and groups in the periodic table?
• Why does atomic radius decrease across a period and increase down a group?
• How does ionization energy change as you move across a period and down a group?

key terms and Concepts

1.7 Seven Electronic Configurations and Orbitals Diagram

The electronic Configuration for any atom follows the following three principles:

• Aufbau principle
• Pauli’s exclusion Principle
• Hund’s rule

Aufbau principle

• In the Ground state of the Atoms, the orbitals are filled in order of their increasing energies.
  • electrons first occupy the lowest energy orbital available to them and enter into higher energy orbitals only after the lower energy orbitals are filled.
• In general, electrons occupy the lowest-energy orbital available before entering the higher energy orbital.
• The relative energy of an orbital is given by (n+l) rule.
• As(n+l)value increases, the energy of orbital increases.
• The orbital with the lowest (n+l)value is filled firsts
• When two or more orbitals have the same (n+l) value, the one with the lowest ‘n’ value (or) highest ‘l ’ value is preferred in filling.

Example

• Consider two orbitals 3d and 4s. n+l value of 3d =3+2=5 and of 4S=4+0=4. Since 4s has lowest (n+l) value, it is filled first before filling taking place in 3d

Example 2, Consider The orbitals 3d,4p and 5s,

• The (n+l) value of 3d =3 +2= 5
• The(n+l) value of4p =4+ 1 =5
• The(n+l) value of 5s =5+ 0=5

These three values are same. Since the‘n’valueis lower to 3d orbitals, the electrons prefer to entering 3d, then 4p and 5s.

• The order of Increasing energy of atomic orbitals Is 1s<2s <2p<3s <3p<4s <3d<4p <5s<4d<5p <6s<4f <5d<6p<7s and so on
• The Sequence In which the Electrons Occupy Various orbitals can be easil remembered with the helpes of Afuba diagrama shown bellow:

Pauli’s Exclusion Principle

• Stated that “No two electrons in an atom can have the same set of values for all the four quantum numbers”
• Two electrons in an orbital may have the same number ,same three of the quantum number but differing spin quantum number.
• In an orbital if one electron has clockwise spin, the other has anticlockwise spin.
• It follows that an orbital can hold a maximum of two electrons with opposite spins.
• One orbital cannot have more than two electron ms
• The maximum capacity holding electron in of amain energy shell is 2n²

Example-

electron configuration N(7) is 1s²s²2p³
The 2p subshell are occupied by a single electron .i.e. 1s² ,2s² ,2p_𝑥¹ ,2p_𝑦¹ ,2p_𝑧¹

Hund’s principle

Hund’s principle: states that electrons will start to fill a set of degenerate orbitals or (equal energy orbitals) keeping the spins parallel  i.e. points up when ms=+ 1/2 and points down when ms=-1/2

Examples: The Ground state of electron structure for the following elements can be represented as:

• Each of the three 2p orbitals(2p_x , 2p_y and 2p_z) will hold a single electron before any of them receives a second electron.

1.7.1 Ground State Electronic Configuration of The Elements

• Two General methods are used to denote electronic configuration

1. The Subshell (sublevel) method

Notation uses numbers to designate the principal energy levels or principal quantum number and the letters s, p, d and f to identify the sublevels or subshell

• Example , hydrogen (H; Z=1) –1s¹
• helium (He; Z=2)– 1s²
• Lithium(Li; Z=3)– 1s²2s¹

2. Orbital diagram method

which consists of a box (or circle, or just a line)for each orbital available in a given energy level, grouped by sublevel, with an arrow indicating the electron’s presence and its direction of spin.

• Example: The Orbital diagrams for the first three elements are

Note: The expected configuration of, chromium and copper based on the aufbau principle, is not the ones observed through the emission spectra and magnetic properties of the elements.

The presence of half-filled and completely filled degenerate orbitals gives greater stability to atoms.

1.8 Electronic Configurations and the Periodic Table of the Elements

1.8.1 The Modern Periodic Table

• The modern periodic law states that the properties of element are periodic functions of their atomic numbers.
• Elements arranged with in the same vertical column, have similar properties and are called Group or Families.
• Elements found in any one of the horizontal row of the periodic table is called period or series.

1.8.2 Classification of Elements

• Elements are arranged in the periodic table in accordance to valence electrons entering the orbital of lowest energy.
• There are 18 groups and 7 periods in the modern periodic table.

• Periodic law states that certain sets of physical and chemical properties recur at regular intervals (periodically) when the elements are arranged according to increasing atomic number.

Representative or main group elements:

• consists of All ”S” and “P” block elements .
• Groups IIIA, IVA, VA, VIA, VII and VIII all have P- orbital filled, and because there properties are dependent on the
  presence of P electrons, they are called the p-block elements.
• The chemical properties of the representative elements are determined by the number of valence electrons in their

The Transition Elements:

• Are elements where d-orbitals are being filled, are also called d-block elements.
• There are four series of transitional elements, 3d, 4d, 5d and 6d depending on the energy level of d-orbitals
• The Inner Transition elements: are elements where the f-orbitals are being filled, are also called f-block elements.
• There are two series of f-block elements, 4f and 5f series called lanthanides and Actinides, respectively

2.8.3 Periodic Properties

1. Atomic Size (Atomic Radii)

1. Are usually obtained by measuring the distance between atoms in a chemical compound.
2. In general, Atomic size increases down the group and decreases across a period of the periodic table.

From left to right, size of atoms decrease in the periodic table.

2. Ionization Energy

• Is the amount of energy required to remove an electron from a gaseous atom or Ion.
  • Eg Na(g) + energy → Na⁺(g) + e
  • In general, Ionization energy decreases down the group and increases across a period of the periodic table.
  • Multi-electron atoms can lose more than one electron, so the ionization energies required to remove each electron are numbered in sequence from the ground-state atom.

• first ionization energy(IE₁) – is the energy needed to remove an electron from the highest occupied sublevel of the gaseous atom.
• The first ionization energy is a key factor in an elements chemical reactivity because, atoms with a low IE₁ tend to form cations during reactions, whereas those with a high IE₁ (except the noble gases) often form anions.

second ionization energy (IE₂) removes the second electron and IE₂ is always larger than IE₁

The Variation In the magnitude of Ionization energy of elements in the periodic table is mainly dependent on the following factors:

a. The size of the atom ( inverse relation IE and atomic size.)

b. The magnitude of the nuclear Charge on the atom,

c. The extent of screening the type of orbital involved(s,p,d, or f), s>p >d>f

In general, ionization energy increases across a period; it is easier to remove an electron from an alkali metal than from a noble gas.

3. The Effective Nuclear charge [ Zeff]:

* Is the actual nuclear charge less the screening effect of other electrons in the atom. Z_eff= Atomic number-core or screening electron
Z_eff = Z – S
Atomic size increases down a group because of an increase in Number of shells
Atomic size decreases across a period due to an increase in Zeff.
Example:- Compare effective Nuclear charges of ₆C, ₇N and ₈O

Solution: Z_eff of C : Z-S= 6-2=+4

Z_eff of N : Z-S= 7-2 =+5

Z_eff of O: Z-S= 8-2=+6
Thus: In an increasing order of Z_eff: C < 𝑁 < O
In an increasing order of size: O < 𝑁 < C

4. Electron Affinity [EA]

 Is the energy change that occurs when an electron is added to a gaseous atom or ion.

• F(g) + e- ⇀ F(g) ……………EA= -328kJ/mol
• Halogen family has the highest electron affinity
• Generally, Electron affinity increases from left to right across a period.
• When multiple electrons are added to a gaseous atom or ion, the separate electron affinities can be written as
• Example:- O(g) + e- ⇀ O^–(g)………………… E_A1= -141KJ/mol
• O^-(g) + e ⇀ O^2-(g) ……………………..EA2 =+744KJ/mol

5. Electronegativity

• Is the power of an atom to attract electrons toward itself when bonded to other atoms.
• On an electronegativity scale devised by Linus Pauling, the most non-metallic element fluorine is assigned the highest value of 4.0.
• Typically active metals have 1.0 or less electro negativities.
• Generally, Electronegativity increases across a period and decreases down a group.

6. Metallic character

* Refer to the chemical properties of metals which arise from their ability to lose electrons.
* Generally, metallic character of elements increases down the group and decreases a cross a period of the periodic table