Periodic Table flash cards

1. 1.State the Modern Periodic Law.

   The physical and chemical properties of elements are a periodic function of their atomic numbers.

   Hint: What is the fundamental property elements are arranged by?

2. 2.Who proposed the Modern Periodic Law and how did it correct Mendeleev's law?

   Henry Moseley. He showed atomic number (not atomic mass) is the fundamental property, resolving anomalies like the Ar/K and Co/Ni pairs.

   Hint: Moseley's work on X-ray spectra.

3. 3.State Mendeleev's Periodic Law.

   The properties of elements are a periodic function of their atomic masses.

   Hint: Older law, based on mass not number.

4. 4.What was Mendeleev's boldest prediction supporting his table?

   He left gaps for undiscovered elements (eka-boron, eka-aluminium, eka-silicon = Sc, Ga, Ge) and predicted their properties accurately.

   Hint: Empty spaces reserved for future elements.

5. 5.State Newlands' Law of Octaves and its limitation.

   Every 8th element has properties similar to the 1st (like musical octaves). It failed beyond calcium.

   Hint: Musical analogy; worked only for lighter elements.

6. 6.What are Dobereiner's Triads?

   Groups of 3 elements where the atomic mass of the middle element is approximately the average of the other two (e.g. Li, Na, K).

   Hint: Sets of three; middle = mean.

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.On what basis is the long form periodic table arranged?

   Based on the electronic configuration of elements; elements in the same group have the same number of valence electrons.

   Hint: Outer-shell electrons decide the group.

9. 9.What does the period number tell you about an element?

   It equals the principal quantum number ($n$) of the outermost (valence) shell being filled.

   Hint: Highest n value = ?

10. 10.How many elements are in each period 1 through 7?

    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.

11. 11.Why does each period start with an alkali metal and end with a noble gas?

    A new period begins when a new principal shell ($n$) starts filling (ns$^1$) and ends when that shell's octet is complete (noble gas configuration).

    Hint: Filling begins ns$^1$, ends with full octet.

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

    A block is a set of elements classified by the subshell (s, p, d, f) into which the last (differentiating) electron enters.

    Hint: Which subshell gets the final electron?

13. 13.Give the general valence-shell configuration of s-block elements.

    $ns^{1-2}$ (Groups 1 and 2: alkali and alkaline earth metals).

    Hint: Groups 1 and 2.

14. 14.Give the general valence-shell configuration of p-block elements.

    $ns^2 np^{1-6}$ (Groups 13 to 18).

    Hint: Groups 13–18, s and p filling.

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

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

    Hint: Penultimate d shell fills.

16. 16.Give the general configuration of f-block (inner transition) elements.

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

    Hint: Antepenultimate f shell fills.

17. 17.Which groups make up the representative (normal) elements?

    The s-block and p-block elements (Groups 1, 2 and 13–18, excluding noble-gas core-filled cases).

    Hint: Main-group elements.

18. 18.Why are d-block elements called transition elements?

    They lie between the s-block (metals) and p-block (non-metals), representing a transition in properties, with partly filled d-orbitals.

    Hint: Bridge between s and p blocks.

19. 19.Where are the f-block elements placed in the periodic table?

    At the bottom, as two separate rows: lanthanoids (58–71) and actinoids (90–103).

    Hint: Two rows shown separately below the main table.

20. 20.To which block do the noble gases belong, and what is the anomaly with He?

    p-block (Group 18), $ns^2np^6$. Helium is $1s^2$ (s-block config) but placed with noble gases due to its properties (full valence shell, inertness).

    Hint: He has no p electrons but sits in Group 18.

21. 21.What are the general characteristics of metals in the periodic table?

    They are on the left and centre; malleable, ductile, good conductors, tend to lose electrons (electropositive).

    Hint: Left side; electropositive.

22. 22.Where are non-metals located and what is their electronic tendency?

    On the upper right side; they tend to gain electrons (electronegative), are poor conductors and often brittle solids/gases.

    Hint: Top-right; electron-accepting.

23. 23.What are metalloids? Give examples.

    Elements along the diagonal staircase (border of metals/non-metals) with intermediate properties: B, Si, Ge, As, Sb, Te.

    Hint: Semi-metals along the zig-zag line.

24. 24.Define effective nuclear charge ($Z_{eff}$).

    The net positive charge experienced by a valence electron after accounting for shielding: $Z_{eff} = Z - S$, where $S$ is the screening constant.

    Hint: Actual pull felt by outer electron = Z minus screening.

25. 25.Define the shielding (screening) effect.

    The reduction of nuclear attraction on outer electrons caused by repulsion from inner-shell electrons that partly block the nuclear charge.

    Hint: Inner electrons shield outer ones from the nucleus.

26. 26.Rank the subshells by their screening ability.

    $s > p > d > f$ (s-electrons shield most effectively, f the least).

    Hint: Order follows penetration power.

27. 27.How does $Z_{eff}$ change across a period and why?

    It increases across a period because nuclear charge rises while electrons are added to the same shell (screening stays roughly constant).

    Hint: More protons, same shell = stronger pull.

28. 28.How does $Z_{eff}$ change down a group?

    It increases only slightly (nearly constant) because added inner shells increase screening, largely offsetting the rise in nuclear charge.

    Hint: Screening cancels most of the added nuclear charge.

29. 29.Define atomic radius and its common types.

    Half the distance between nuclei of adjacent bonded atoms. Types: covalent radius, metallic radius, and van der Waals radius (for noble gases).

    Hint: Half the internuclear distance; three kinds.

30. 30.Why is van der Waals radius larger than covalent radius?

    It is half the distance between non-bonded atoms (only weak forces), so nuclei are farther apart than in a covalent bond.

    Hint: Non-bonded contact vs bonded overlap.

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