Electronic Configuration of Elements - UNSOLVED PRACTICE SET

Q7. Write the electronic configuration of the following elements using the noble gas core notation:

(a) Potassium (Z = 19)

(b) Bromine (Z = 35)

Q8. What is meant by valence electrons? How do valence electrons determine the chemical properties of an element? 

Q9. Write the electronic configuration of Fe²⁺ and Fe³⁺ ions. Which ion is more stable and why? 

Q10. How does the electronic configuration determine whether an element belongs to the s-block, p-block, d-block, or f-block?

Q11. Your chemistry teacher explains that sodium (Na) is highly reactive because it has one electron in its outermost shell, while neon (Ne) is inert because its outermost shell is full. Using electronic configurations, explain why potassium (K) is even more reactive than sodium, even though both have one valence electron.

Q12. Write the electronic configuration of copper (Z = 29). Why does it deviate from the expected Aufbau configuration?

Q13. Explain how the electronic configuration of elements is written using the Aufbau principle, Pauli exclusion principle, and Hund's rule. Write the complete electronic configurations (in both standard and noble gas notation) for:

(a) Sulphur (Z = 16)

(b) Zinc (Z = 30)

(c) Krypton (Z = 36)

Also identify the block to which each element belongs.

Q14. Describe the relationship between electronic configuration and the position of an element in the periodic table. Explain how the configuration determines:

(a) The period number

(b) The group number

(c) The block (s, p, d, or f)

For an element with configuration [Ar] 3d⁵ 4s², determine its period, group, block, and name.

Q15. During a chemistry practical exam, you are given four unknown elements with the following properties:

Element A: Forms a +1 ion by losing one electron; its ion has the configuration of neon.

Element B: Has 6 valence electrons and forms a −2 ion.

Element C: Has the configuration [Ar] 3d³ 4s² and forms coloured compounds.

Element D: Is a noble gas with atomic number 10.

(a) Identify each element and write its electronic configuration.

(b) For Element A, write the configuration of its +1 ion. Is the ion larger or smaller than the neutral atom? Explain.

(c) For Element C, predict how many unpaired electrons it has in its ground state. Is it paramagnetic?

(d) Element E has atomic number 24. Write its expected and actual configurations. Explain the stability of its actual configuration using the concepts of half-filled subshells.

 Q16. The following table gives information about certain elements:

(a) Complete the table by writing the electronic configuration of each element.

(b) Determine the group, period, and block for each element.

(c) Which element would have the highest ionization energy? Explain using electronic configuration.

(d) Which two elements would form an ionic compound with each other? Write the formula of the compound.

Q17. A transition metal M has the electronic configuration [Ar] 3d⁷ 4s² in its ground state.

(a) Identify the element M.

(b) Write the electronic configurations of M²⁺ and M³⁺ ions.

(c) Which ion — M²⁺ or M³⁺ — would have more unpaired electrons? Show your reasoning with orbital diagrams.

(d) Predict the magnetic behaviour (paramagnetic/diamagnetic) of M, M²⁺, and M³⁺. Explain your predictions.

(e) The element just before M in the periodic table has configuration [Ar] 3d⁵ 4s¹. Explain why this configuration is more stable than [Ar] 3d⁴ 4s².

Q18. Write the electronic configurations of the following species and determine the number of unpaired electrons in each:

(a) Nitrogen atom (Z = 7)

(b) Oxygen atom (Z = 8)

(c) Cr³⁺ ion (Z = 24)

(d) Cu⁺ ion (Z = 29)

For each, draw the orbital diagram for the valence shell.