Classification of Elements and Periodicity in Properties

The periodic table developed through Dobereiner's triads, in which the middle element's atomic mass is about the average of the other two (Li 7, Na 23, K 39); Chancourtois's arrangement; Newlands's law of octaves; and Lothar Meyer's curves. It then covers Mendeleev's periodic law of 1869. Mendeleev left gaps and predicted the properties of eka-aluminium (gallium) and eka-silicon (germanium). Moseley's work on X-ray spectra showed that atomic number is the more fundamental property, leading to the modern periodic law and the long form of the periodic table with 18 groups and 7 periods. You will learn the IUPAC nomenclature for elements with atomic numbers above 100, such as unbinilium for Z = 120, and how electronic configurations place elements in the s, p, d and f blocks. The chapter explains periodic trends in physical and chemical properties: atomic and ionic radii, isoelectronic species, ionisation enthalpy, electron gain enthalpy, electronegativity and valence. It also explains the anomalous behaviour of second-period elements, diagonal relationships and the nature of oxides.

Detailed Worksheet: Classification of Elements and Periodicity in Properties Section A - Definitions (10 marks) 1. State the modern periodic law. (2 marks) 2. What are Dobereiner's triads? Give one example. (2 marks) 3. Define ionisation enthalpy. Why is the second ionisation enthalpy higher than the first? (2 marks) 4. What are isoelectronic species? Give two examples. (2 marks) 5. Define electronegativity. Which element is the most electronegative? (2 marks) Section B - Calculations and Applications (15 marks) 6. Using Dobereiner's triads, calculate the average atomic mass of the first and third elements and compare it with the middle one: (i) Ca 40, Sr 88, Ba 137 (ii) Cl 35.5, Br 80, I 127. Comment on the results. (3 marks) 7. Write the IUPAC names and symbols of elements with atomic numbers 101, 109, 115 and 120, using the numerical roots nil (0), un (1), bi (2), tri (3), quad (4), pent (5), hex (6), sept (7), oct (8), enn (9). (3 marks) 8. The first ionisation enthalpy of sodium is 496 kJ/mol. Calculate the energy needed to ionise 2.3 g of sodium vapour completely to Na+ ions, and the energy needed per atom in joules (Na = 23, NA = 6.022 x 10^23). (3 marks) 9. Write the electronic configuration and identify the period, group and block of elements with atomic numbers 17, 20, 26 and 35. (3 marks) 10. Arrange in order of increasing ionic radius and explain: N3-, O2-, F-, Na+, Mg2+, Al3+. What do these ions have in common? (3 marks) Section C - Diagrams (10 marks) 11. Draw an outline of the long form of the periodic table, marking the s, p, d and f blocks with group numbers. (4 marks) 12. Draw a graph showing the variation of first ionisation enthalpy with atomic number for elements Z = 1 to 20. (3 marks) 13. Draw a diagram showing the trends of atomic radius, ionisation enthalpy, electron gain enthalpy and electronegativity across a period and down a group. (3 marks) Section D - Analysis and Higher-order Thinking (15 marks) 14. Why does the first ionisation enthalpy generally increase across a period? Explain why boron has a lower first ionisation enthalpy than beryllium, and oxygen lower than nitrogen. (5 marks) 15. Explain why the electron gain enthalpy of fluorine is less negative than that of chlorine, and why noble gases have large positive electron gain enthalpies. (5 marks) 16. Explain the anomalous behaviour of second-period elements and the diagonal relationship between lithium and magnesium. How do periodic trends determine the nature of oxides across period 3? (5 marks) Instructions: Time allowed 2 hours. Attempt all sections. Show the steps of all calculations with units. Draw graphs and tables neatly.