1.2.1 First and successive ionisation energies Quiz: Pearson Edexcel Chemistry, Unit 1

The 20 questions

1. Define first ionisation energy.

   • The energy needed to remove one electron from each atom in one mole of gaseous atoms
   • The energy needed to form one mole of an ionic compound from its elements in their standard states
   • The energy released when one mole of gaseous atoms gains an electron to form negative ions
   • The energy needed to remove one electron from a single gaseous atom, measured in joules per atom

2. Which equation represents the first ionisation energy of sodium?

   • Na(g) -> Na+(g) + e-
   • Na(g) + e- -> Na-(g)
   • Na+(g) -> Na2+(g) + e-
   • Na(s) -> Na+(s) + e-

3. Which equation represents the second ionisation energy of magnesium?

   • Mg(g) -> Mg2+(g) + 2e-
   • Mg2+(g) -> Mg3+(g) + e-
   • Mg(s) -> Mg+(s) + e-
   • Mg+(g) -> Mg2+(g) + e-

4. Why are successive ionisation energies always larger than the one before?

   • The nuclear charge decreases after each electron is removed
   • Electron-electron repulsion increases, pulling electrons away from the nucleus
   • Each electron removed is closer to the nucleus of a neutral atom
   • Each electron is removed from a more positive ion, so it is held more strongly

5. Which factor does NOT directly increase first ionisation energy?

   • Increasing the number of protons in the nucleus
   • Reducing the shielding by inner electrons
   • Removing the electron from a smaller-radius orbital
   • Increasing the number of neutrons in the nucleus

6. Which element in Period 3 has the lowest first ionisation energy?

   • Magnesium
   • Sodium
   • Argon
   • Aluminium

7. Successive ionisation energies (kJ/mol) of an element are 578, 1817, 2745, 11578. How many electrons are in its outer shell?

   • 4
   • 3
   • 2
   • 1

8. Successive ionisation energies (kJ/mol) of an element X are 738, 1451, 7733, 10540. Which group does X belong to?

   • Group 1
   • Group 2
   • Group 4
   • Group 3

9. Successive ionisation energies (kJ/mol) of X are 496, 4562, 6910, 9543. What is the most likely group of X?

   • Group 2
   • Group 1
   • Group 3
   • Group 7

10. Define successive ionisation energy.

    • The energy needed to remove all of the electrons from one mole of a solid element at room temperature
    • The energy needed to add an electron to each atom in one mole of gaseous atoms of an element
    • The energy released when one mole of gaseous ions forms a covalent bond with another ion
    • The energy needed to remove one electron from each ion in one mole of gaseous ions with a given positive charge

11. Why is the first ionisation energy of boron lower than that of beryllium?

    • Boron's outer electron is in a lower-energy 1s orbital, which is less shielded than beryllium's outer electron in the atom
    • Boron's 2p sub-shell is full, so its outer electron resists removal more than beryllium's electron does in the atom
    • Boron has fewer protons than beryllium, so its nucleus attracts the outer electrons less strongly, which lowers the energy needed
    • Boron's outer electron is in a 2p sub-shell, which is higher in energy and more shielded than the 2s electrons of beryllium

12. Why does the first ionisation energy of Li exceed that of Na?

    • Sodium's outer electron is in a lower-energy sub-shell, so it is harder to remove from the atom
    • Lithium's outer electron is in the second shell, closer to the nucleus with less shielding
    • Sodium has a larger nuclear charge per electron than lithium, which pulls its outer electron closer
    • Lithium has more neutrons than sodium, so its nucleus holds the outer electron more firmly

13. Which statement describes the trend in first ionisation energy across Period 2?

    • It generally increases from lithium to neon, with small dips at boron and oxygen
    • It stays almost constant apart from a single large jump at neon, which has a full outer shell
    • It decreases steadily from lithium to neon, because each extra electron is held less strongly
    • It increases only from beryllium to nitrogen, and then falls back to its starting value at neon

14. Which element has the greatest first ionisation energy?

    • Fluorine
    • Neon
    • Oxygen
    • Sodium

15. Explain why the first ionisation energy of oxygen is lower than that of nitrogen.

    • Oxygen's 2p electron is further from the nucleus than nitrogen's, so it is removed more easily
    • Oxygen's outer electrons are all in a completely filled sub-shell, which weakens their attraction
    • Oxygen's removed electron comes from a paired 2p orbital, and repulsion within the pair makes it easier to remove
    • Oxygen has fewer protons than nitrogen, so its nucleus attracts the outer electrons less strongly overall

16. A graph of successive ionisation energies for an element shows a large jump between the fourth and fifth values. Which best explains this?

    • The fifth electron is in the same shell as the others but has a greater spin, so it is held more weakly
    • The fifth electron is removed from the outer shell, which has more shielding from the inner electrons
    • The fifth electron is removed from a new inner quantum shell, closer to the nucleus, so it is held much more strongly
    • The nuclear charge drops sharply after four electrons are removed, reducing the attraction on the rest

17. Explain why the second ionisation energy of an element is greater than its first.

    • The nucleus loses protons after the first electron is removed, which lowers the attraction on the rest
    • Electron pairing lowers the energy needed to remove the second electron from the positive ion
    • The second electron is in a higher-energy shell than the first electron, so it is further from the nucleus
    • Removing an electron from a positive ion increases the attraction between the nucleus and the remaining electrons

18. Explain why the first ionisation energy of aluminium is lower than that of magnesium.

    • Aluminium's outer electron is in a 3p sub-shell, which is higher in energy and more shielded than the 3s electrons of magnesium
    • Aluminium has fewer protons than magnesium, so its nucleus attracts its outer electron less strongly overall
    • Aluminium's outer electron is in a 2p sub-shell, which is closer to the nucleus than the 3s electrons
    • Magnesium's outer electron is in a higher-energy sub-shell than aluminium's outer electron, so it is easier to remove

19. The first ionisation energy of potassium is lower than that of lithium. Which factor is most important?

    • Potassium has fewer protons than lithium, so its nucleus attracts the outer electron more weakly
    • Potassium's outer electron is in the same shell as lithium's, so the shielding is identical in both
    • Potassium has more inner shells, so its outer electron is further from the nucleus and more shielded
    • Potassium has a higher electronegativity than lithium, so its outer electron is pulled away more easily

20. Which factor increases the first ionisation energy of an element most directly?

    • More neutrons in the nucleus, which add to the mass
    • A larger atomic radius for the outer electron
    • Greater shielding from inner electron shells
    • A greater nuclear charge with similar shielding

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