Periodic Table & Periodicity flash cards

1. 1.State the Modern Periodic Law.

   Physical and chemical properties of elements are a periodic function of their atomic numbers ($Z$), not atomic masses.

   Hint: What did Moseley change?

2. 2.State Mendeleev's Periodic Law.

   Properties of elements are a periodic function of their atomic masses.

   Hint: Basis of arrangement in 1869.

3. 3.What experiment led to atomic number replacing atomic mass as the basis of periodicity?

   Moseley's X-ray studies (1913): frequency of characteristic X-rays gave $\sqrt{\nu} \propto (Z-b)$, showing $Z$ is the fundamental property.

   Hint: Anode metals and X-ray lines.

4. 4.Name two major achievements of Mendeleev's table.

   (1) Left gaps for undiscovered elements (predicted eka-boron=Sc, eka-aluminium=Ga, eka-silicon=Ge). (2) Corrected doubtful atomic masses (e.g. Be, In).

   Hint: Predictions and corrections.

5. 5.List three key defects of Mendeleev's periodic table.

   (1) No fixed place for hydrogen. (2) Isotopes had no place. (3) Some pairs placed in anomalous order by mass (e.g. Ar before K, Co before Ni, Te before I).

   Hint: H, isotopes, mass anomalies.

6. 6.Why is Ar (mass 39.9) placed before K (mass 39.1) in the modern table though it violates Mendeleev's mass order?

   Modern table orders by atomic number: $Z_{Ar}=18 < Z_{K}=19$. The anomaly vanishes.

   Hint: Ordering property matters.

7. 7.How many periods and groups are in the modern long-form periodic table?

   7 periods (horizontal rows) and 18 groups (vertical columns).

   Hint: Rows vs columns count.

8. 8.What does the period number of an element indicate?

   The principal quantum number ($n$) of its outermost (valence) shell.

   Hint: Highest $n$ occupied.

9. 9.Give the number of elements in each of the 7 periods.

   Period 1: 2; Period 2: 8; Period 3: 8; Period 4: 18; Period 5: 18; Period 6: 32; Period 7: 32.

   Hint: 2, 8, 8, 18, 18, 32, 32.

10. 10.Why does Period 1 contain only 2 elements?

    Only the $1s$ orbital is available ($n=1$ has just one orbital), holding a maximum of 2 electrons (H, He).

    Hint: Orbitals available at $n=1$.

11. 11.The number of elements in a period equals the number of electrons in which set of orbitals?

    The set filled at that energy level; period length = number of electrons accommodated between two successive noble gases (governed by $(n-1)d$, $(n-2)f$, $ns$, $np$ filling order).

    Hint: Aufbau across a period.

12. 12.Define a block in the periodic table.

    A set of elements grouped by the subshell that receives the last (differentiating) electron: s, p, d, or f block.

    Hint: Last electron enters which subshell.

13. 13.Which groups constitute the s-block, and what is their general valence configuration?

    Groups 1 and 2 (alkali & alkaline earth metals), plus He. Configuration $ns^{1-2}$.

    Hint: Two leftmost columns.

14. 14.Which groups form the p-block and what is the valence configuration?

    Groups 13–18. Configuration $ns^2\,np^{1-6}$ (He is $1s^2$ but grouped with 18).

    Hint: Right side, includes metalloids & noble gases.

15. 15.Give the general valence configuration of d-block (transition) elements.

    $(n-1)d^{1-10}\,ns^{0-2}$ — groups 3 to 12.

    Hint: Penultimate d subshell filling.

16. 16.Give the general valence configuration of f-block elements and name the two series.

    $(n-2)f^{1-14}\,(n-1)d^{0-1}\,ns^2$; the lanthanoids ($4f$) and actinoids ($5f$).

    Hint: Inner transition, antepenultimate f.

17. 17.Why is He placed in Group 18 (p-block position) despite an $s^2$ configuration?

    Its properties (fully filled, chemically inert, noble gas) match Group 18, not the reactive s-block metals of Group 2.

    Hint: Property over configuration.

18. 18.Why are transition elements so named?

    They lie between the s-block (metals) and p-block (non-metals) and represent a transition in properties; d-electrons fill the penultimate shell.

    Hint: Bridge between blocks.

19. 19.Which elements are called typical or representative elements?

    The s-block and p-block elements (main-group, excluding noble gases sometimes).

    Hint: Not d or f block.

20. 20.Define effective nuclear charge $Z_{eff}$.

    The net positive charge actually experienced by a valence electron: $Z_{eff}=Z-\sigma$, where $Z$ is nuclear charge and $\sigma$ is the screening (shielding) constant.

    Hint: Actual pull felt by outer electron.

21. 21.What is the screening (shielding) effect?

    The reduction of nuclear attraction on outer electrons due to repulsion by inner-shell electrons, which partly cancel the nuclear charge.

    Hint: Inner electrons blocking the nucleus.

22. 22.Order the screening ability of orbital types for a given shell.

    $s > p > d > f$. s-electrons shield (and penetrate) best; f-electrons shield poorly.

    Hint: Penetration order.

23. 23.How does $Z_{eff}$ vary across a period, and why?

    $Z_{eff}$ increases across a period: nuclear charge rises while electrons enter the same shell, giving poor mutual shielding.

    Hint: Same shell, more protons.

24. 24.How does $Z_{eff}$ (felt by valence electrons) change down a group?

    It increases only slightly; added inner shells shield the added nuclear charge quite effectively, so outer electrons feel a nearly constant pull.

    Hint: New shells shield well.

25. 25.Define atomic radius and name its common types.

    Half the internuclear distance between bonded atoms. Types: covalent radius, metallic radius, and van der Waals radius (for non-bonded contact).

    Hint: Half the bond length.

26. 26.Rank covalent, metallic, and van der Waals radii for a given element.

    $r_{covalent} < r_{metallic} < r_{vanderWaals}$.

    Hint: vdW is largest (weak non-bonded).

27. 27.Why do noble gases show anomalously large atomic radii?

    Their radius is measured as a van der Waals radius (non-bonded), which is inherently larger than covalent/metallic radii of neighbours.

    Hint: Different radius type measured.

28. 28.How does atomic radius vary across a period? Explain.

    It decreases left to right: increasing $Z_{eff}$ pulls the same-shell electrons closer.

    Hint: $Z_{eff}$ rising, shell fixed.

29. 29.How does atomic radius vary down a group? Explain.

    It increases: each period adds a new principal shell, increasing $n$ and outweighing the rise in nuclear charge.

    Hint: New shell each period.

30. 30.Define ionic radius. How does a cation compare with its parent atom?

    Radius of an ion in an ionic crystal. A cation is smaller than its parent atom (fewer electrons, often loses a shell, higher $Z_{eff}$ per electron).

    Hint: Loss of electrons shrinks.

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