Electron Configuration
An electron configuration is a description of electron arrangement within an atom, which indicates both population and location of electrons among the various atomic orbitals.

General Rules for Electron Configurations
1.   Electrons occupy orbitals of the lowest energy available.

2.   There can be a maximum of only two electrons in any given orbital.  No two electrons in the same atom may have all four quantum numbers alike. This is a statement of the Pauli Exclusion Principle.

3.     Building-Up Principle (Aufbau Principle):  Electrons are added by successively filling subshells with electrons in a specific order based on increasing energies of the subshells.
        The maximum number of electrons in any s subshell is two.
        The maximum number of electrons in any p subshell is six.
        The maximum number of electrons in any d subshell is ten.
        The maximum number of electrons in any f subshell is fourteen.

4.   Ground State:  The electron configuration associated with the lowest energy level of the atom is referred to as ground state. Each electron in the atom or ion will be in the lowest energy level possible. Configurations associated with electrons in energy levels other than the lowest are referred to as excited states.

Example Configurations
Hydrogen has a single electron and therefore has the following configuration.   H 1s¹

    The numeral 1 refers to the value of n, the principal quantum number.

    The letter s refers to the l value, the angular momentum number.  Recall that when l = 0 there
    is a letter designation of s.

    The numeral 1 in the superscript refers to the number of electrons in the 1s subshell.

Nitrogen has 7 electrons which are distributed as shown in the following configuration.  N 1s² 2s² 2p³

Core Electrons (Noble Gas Core)
The terms "core electrons" or "noble gas core" refer to the electrons within the atom which have the
same electron configuration as the nearest noble gas of lower atomic number.  The core electrons are
the inner electrons which are not directly involved in bonding.

The core for Li is He.  The core electrons of Li have the identical electron configuration as an atom of He.
The core for F is also He.  The core for Al is Ne.

The electron configuration for argon is :  Ar 1s² 2s² 2p⁶ 3s² 3p⁶

The electron configuration for potassium is:  K 1s² 2s² 2p⁶ 3s² 3p⁶ 4s¹

Potassium has an argon core plus 4s¹

An abbreviated method for electron configurations uses a set of square brackets [ ] around the chemical symbol of the noble gas.

The abbreviated electron configuration for potassium is: K [Ar] 4s¹

Chlorine has a neon core.

The configuration for chlorine is:  Cl 1s² 2s² 2p⁶ 3s² 3p⁵ or the abbreviated method is:  Cl [Ne] 3s² 3p⁵

Then noble gas core together with (n-1)d¹⁰ is known as the pseudo-noble-gas core.

Valence Electrons
Valence electrons are those electrons in an atom outside the noble-gas core.  Valance electrons are the electrons in the outermost principal quantum level.  Valance electrons are the electrons which are most often involved in reactions and forming chemical bonds.

Potassium has a single valence electron, 4s¹, which comes from the 4s subshell.

Chlorine has a total of seven valence electrons, 3s² 3p⁵, two from the 3s subshell and five from the 3p subshell.

Electron Configurations and the Periodic Table
The representative elements (also called main group elements) are the elements in Groups 1 (1A)  through 17 (7A), all of which have incompletely filled s or p subshells of the highest principal quantum number. The representative elements all have valence shell configurations of ns^anp^b, with some choice of a and b.

s-block elements
Group 1 (1A; the alkali metals) and Group 2 (2A; the alkaline earth metals) are referred to as s-block elements

Group 1 elements have a noble gas core plus 1 valence electron with an ns¹ configuration.

    Li [He] 2s¹   or  K [Ar] 4s¹

Group 2 elements have a noble gas core plus 2 valence electrons with an ns² configuration

    Be [He] 2s²  or  Ca [Ar] 4s²

p-block elements
Group 13 (3A) elements through Group 18  (8A) are referred to as p-block elements.

Group 13 elements have the general configuration of ns² np¹
    B [He] 2s² 2p¹

Group 14 elements have the general configuration of ns² np²
    C [He] 2s² 2p²

Group 15 elements have the general configuration of ns² np³
    N [He] 2s² 2p³

Group 16 elements have the general configuration of ns² np⁴
    O [He] 2s² 2p⁴

Group 17 elements have the general configuration of ns² np⁵
    F [He] 2s² 2p⁵

The p-block elements in the fourth period and beyond will have the noble gas core together with (n-1)d¹⁰.
    Br [Ar] 4s² 3d¹⁰ 4p⁵

In Group 18 (8A; the noble gases) the p subshell has just been filled.

    Ar 1s² 2s² 2p⁶ 3s² 3p⁶

Transition Elements, the d-block elements
In the d-block transition elements (transition metals) a d subshell is being filled. The principal quantum
number of the d subshell is always 1 less than the period in which the element is located. Technetium
(Tc, atomic number = 43) is in the fifth period.

    Tc [Kr] 5s² 4d⁵

Inner Transition Elements, the f-block elements
In the f-block transition elements (inner transition) an f subshell is being filled. The principal quantum
number of the f subshell is always 2 less than the period in which the element is located. Plutonium
(Pu, atomic number = 94) is in the seventh period.  The last electrons placed in the electron configuration
go into the 5f subshell.

There are a few exceptions to the predicted electron configurations. The predicted and observed
configurations for some of the exceptions are shown below.
    Predicted                      Observed
    Cr [Ar] 4s² 3d⁴             Cr [Ar] 4s¹ 3d⁵

    Cu [Ar] 4s² 3d⁹            Cu [Ar] 4s¹ 3d¹⁰

    Ag [Kr] 5s² 4d⁹            Ag [Kr] 5s¹ 4d¹⁰

    N: 1s² 2s² 2p³             N^3-: 1s² 2s² 2p⁶

    O: 1s² 2s² 2p⁴            O^2-: 1s² 2s² 2p⁶

Electron Configurations of Cations
To write the ground state electron configuration of a cation, remove electrons from the highest occupied energy level in the ground state electron configuration of the atom.   In other words, remove electrons from the orbital with the highest principal quantum number.

Na: 1s² 2s² 2p⁶ 3s¹         Na⁺: 1s² 2s² 2p⁶

Mg: 1s² 2s² 2p⁶ 3s²        Mg²⁺: 1s² 2s² 2p⁶

Fe: 1s² 2s² 2p⁶ 3s² 3p⁶ 4s² 3d⁶

Fe²⁺: 1s² 2s² 2p⁶ 3s² 3p⁶ 3d⁶

Fe³⁺: 1s² 2s² 2p⁶ 3s² 3p⁶ 3d⁵

Isoelectronic  Atoms & ions that possess the same number of electrons, and hence the same ground state electron configuration are said to be isoelectronic.   Ne, O^2-, Na⁺ are isoelectronic, which is to say that these three have identical electron configurations.

Hydrogen has a single electron in a 1s orbital.  The orbital diagram for hydrogen is:
H  C
      1s

Helium has two electrons in a 1s orbital.  The orbital diagram is:
He CD
      1s

Other examples of orbital diagrams are:

Li  CD  C
       1s        2s
 

B CD  CD  C  ____  ____
      1s        2s                    2p

Hund's Rule
Hund's rule states that the lowest energy arrangement of electrons in a subshell is obtained by putting electrons into separate orbitals of the subshell with the same spin (parallel spins) before pairing electrons.   When putting electrons into orbitals of the same energy, one electron will occupy each orbital before any orbital has two electrons.   Unpaired electrons have parallel spins.   Paired electrons have unparallel or opposite spins.   This rule is in agreement with the idea that electrons stay as far apart as possible.   This arrangement of electrons makes the total energy of an atom as low as possible.

C CD   CD  C  C  ____
      1s          2s                      2p

N CD   CD  C  C  C .
      1s             2s                      2p

O CD   CD  CD  C  C
      1s           2s                2p

F CD  CD  CD CD  C
     1s        2s                 2p

Ne CD  CD  CD  CD CD
      1s         2s                  2p

Atoms with unpaired electrons are said to be paramagnetic. These are weakly attracted to a magnetic field.

Atoms with all paired electrons are said to be diamagnetic. These are weakly repelled from a magnetic field.

Atomic radius increases right to left and top to bottom .
An atom does not have a definite size, because the statistical distribution of electrons does not
abruptly end but merely decreases to very small values as the distance from the nucleus increases.

1. Within each period, the atomic radius tends to decrease with increasing atomic number.
        K > Ca > Sc
2. Within each group, the atomic radius tends to increase with the period number.
        Li < Na < K

Ionic Radius
The order of ionic radii size is: Cation < Atom < Anion  (Na¹⁺ < Na < Na^1-)

Ionization energy increases left to right and bottom to top.
Ionization energy (IE) is the minimum energy needed to remove the highest energy (outermost) electron
from the neutral atom in the gaseous state.  IE₁ is the energy required to remove the first electron.
IE₂ is the energy required to remove the second electron.

Li(g)  ----> Li¹⁺(g) + e^-   IE₁ = +520 kJ/mol
Na(g)  ----> Na¹⁺(g) + e^-   IE₁ = +496 kJ/mol
K(g)  ----> K¹⁺(g) + e^-   IE₁ = +419 kJ/mol

Li(g)  ----> Li¹⁺(g) + e^-   IE₁ = +520 kJ/mol
Be(g)  ----> Be¹⁺(g) + e^-   IE₁ = +899 kJ/mol

1. Draw orbital diagrams for the valence electrons for each of the following. Which would exhibit paramagnetism?
      a) C    b) O    c) N^3-    d) Mn²⁺    e) Sc³⁺

2. Give full and abbreviated electron configurations for:    a) Br    b) Ag    c) Fe    d) S^2-    e) Ni²⁺

3. For each of the following pairs of orbitals, indicate which orbital is higher in energy:
      a) 1s, 2s     b) 2p, 3p     c) 3d_xy, 3d_yz     d) 4s, 3d     e) 5s, 4f

4. Arrange the following atoms in order of increasing size:    Ar, Ca, K, Sc.

5. Arrange the following in order of decreasing size:    Br^-, Cl^-, F^-, I^-.

6. Indicate which of the elements are s-block, p-block, alkali metals, etc. as well as metal or nonmetal.
      a) Sc     b) P     c) Pu     d) Fr     e) Ni⁺²     f) As

7. What is the symbol for the second period Group IVA element?

8. What group does the element belong to whose X²⁺ ion has 1 unpaired electron in its ground state?  IIIA, IVA or IB?

9. Would element 117 more likely form a stable anion or stable cation and what would its charge be?

10. Would element 119 form a stable anion or stable cation and what would its charge be?

11. How many different neutral species can have the following configurations?           a) ns² np²      b) ns² n-1d⁵

Answers:
1.   a) (paramagnetic)    b) (paramagnetic)     c) (diamagnetic)   d) (paramagnetic)

2.   a) 1s² 2s² 2p⁶ 3s² 3p⁶ 4s² 3d¹⁰ 4p⁵, or [Ar] 4s² 3d¹⁰ 4p⁵
      b) 1s² 2s² 2p⁶ 3s² 3p⁶ 4s² 3d¹⁰ 4p⁶ 5s¹ 4d¹⁰, or [Kr] 5s¹ 4d¹⁰
      c) 1s² 2s² 2p⁶ 3s² 3p⁶ 4s² 3d⁶, or [Ar] 4s² 3d⁶
      d) 1s² 2s² 2p⁶ 3s² 3p⁶, or [Ar] or [Ne] 3s² 3p⁶
      e) 1s² 2s² 2p⁶ 3s² 3p⁶ 4s⁰ 3d⁸, or [Ar] 3d⁸

3.    a) 2s      b) 3p      c) equal      d) 3d      e) 4f

4.       Ar < Sc < Ca < K

5.       F < Cl < Br < I

6.      a) d-block element or transition metal       b) p-block element or nonmetal       c) f-block element or actinide
         d) s-block element or alkali metal        e) d-block element or transition metal        f) p-block element or metalloid

7.      C

8.      IIIA

9.      Anion with a -1 charge.

10.      Cation with a +1 charge.

11.      a)  C & Si        b)  Mn & Tc