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CHEMISTRY, M1 EQ-Bank 14

Compare and explain the reactivity of Group 1 (alkali metals) and Group 2 (alkaline earth metals) with water. In your answer, link your explanation to electron configuration, atomic radius, and ionisation energy.   (6 marks)

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  • Group 1 metals (e.g. \(\ce{Li, Na, K}\)) react vigorously with water to form a hydroxide and hydrogen gas.
  • Group 2 metals (e.g. \(\ce{Mg, Ca}\)) also react with water but much less vigorously, especially at the top of the group. For example, magnesium reacts only slowly with cold water.
  • Electron configuration: Group 1 metals have one valence electron, while Group 2 metals have two valence electrons. Losing one electron requires less energy than losing two, making Group 1 metals more reactive.
  • Atomic radius and ionisation energy: Down both groups, the atomic radius increases, shielding increases, and ionisation energy decreases. This means reactivity with water increases down the group.
  • Therefore: Reactivity increases down both groups, but Group 1 metals show higher reactivity with water compared with Group 2 metals in the same period.
Show Worked Solution
  • Group 1 metals (e.g. \(\ce{Li, Na, K}\)) react vigorously with water to form a hydroxide and hydrogen gas.
  • Group 2 metals (e.g. \(\ce{Mg, Ca}\)) also react with water but much less vigorously, especially at the top of the group. For example, magnesium reacts only slowly with cold water.
  • Electron configuration: Group 1 metals have one valence electron, while Group 2 metals have two valence electrons. Losing one electron requires less energy than losing two, making Group 1 metals more reactive.
  • Atomic radius and ionisation energy: Down both groups, the atomic radius increases, shielding increases, and ionisation energy decreases. This means reactivity with water increases down the group.
  • Therefore: Reactivity increases down both groups, but Group 1 metals show higher reactivity with water compared with Group 2 metals in the same period.

CHEMISTRY, M1 EQ-Bank 5 MC

Which of the following elements has the highest second ionisation energy?

  1. Sodium
  2. Magnesium
  3. Aluminium
  4. Silicon
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\(A\)

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  • Second ionisation energy = energy required to remove a second electron after the first has already been removed.
  • Sodium: After losing one electron, it becomes \(\ce{Na^+}\) with a stable noble gas configuration (like neon). Removing a second electron means breaking into the stable core → very high second IE.
  • The second electron still comes from the valence electron shell for the other three elements hence they all have relatively low second ionisation energies.
  • Therefore, the element with the highest second ionisation energy is Sodium.

\(\Rightarrow A\)

CHEMISTRY, M1 EQ-Bank 2 MC

Ionisation energy is the energy required to remove an electron from an atom.

The table below shows the ionisations energies for element \(\text{Y}\). 

\begin{array} {|l|c|c|}
\hline \text{Ionisation energy number} & 1\text{st} & 2\text{nd} & 3\text{rd} & 4\text{th} & 5\text{th} & 6\text{th}\\
\hline \text{Ionisation energy (kJ/mol)} & 738 & 1450 & 7730 & 10\,500 & 13\,600 & 18\,000 \\
\hline \end{array}

Identify the element described by the table.

  1. Sodium
  2. Magnesium
  3. Silicon
  4. Phosphorus
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\(B\)

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  • The big jump occurs between the 2nd (1450) and 3rd (7730) ionisation energies.
  • This means the element has two valence electrons (after those are removed, you hit a stable core).
  • Group two elements have two valence electrons. 

\(\Rightarrow B\)

CHEMISTRY, M1 EQ-Bank 1 MC

Consider the elements in Group 1 of the periodic table.

What is the general trend in the first ionisation energy of these elements as you move down the group?

  1. Increases from lithium to caesium
  2. Decreases from lithium to caesium
  3. Increases from lithium to potassium then decreases to caesium
  4. Remains constant from lithium to caesium
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\(B\)

Show Worked Solution
  • The first ionisation energy decreases down the group, this is because the outermost electron is further from the nucleus and more shielded by inner electron shells, making it easier to remove.

\(\Rightarrow B\)

CHEMISTRY, M1 EQ-Bank 8

Using your understanding of periodic trends, explain and predict the differences in the properties of the elements lithium \(\ce{(Li)}\) and fluorine \(\ce{(F)}\) regarding their:

    • atomic radii
    • first ionisation energy
    • electronegativity

Give reasons for your predictions based on their positions in the periodic table and electronic configurations.   (4 marks)

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Atomic Radii:

  • Lithium, being in the same period but a different group than fluorine, has fewer protons and a less effective nuclear charge, resulting in a larger atomic radius.

First Ionisation Energy:

  • Fluorine has a higher first ionisation energy due to its greater nuclear charge and smaller radius, which strongly attracts electrons.

Electronegativity:

  • Fluorine, being one of the most electronegative elements, has a strong ability to attract bonding electrons.
Show Worked Solution

Atomic Radii:

  • Lithium, being in the same period but a different group than fluorine, has fewer protons and a less effective nuclear charge, resulting in a larger atomic radius.

First Ionisation Energy:

  • Fluorine has a higher first ionisation energy due to its greater nuclear charge and smaller radius, which strongly attracts electrons.

Electronegativity:

  • Fluorine, being one of the most electronegative elements, has a strong ability to attract bonding electrons.

CHEMISTRY, M1 2009 HSC 30c

The graph shows the first ionisation energy of some elements.
 

Account for the trends in the graph in terms of the electron configuration of the elements.  (3 marks)

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Show Worked Solution
  • The increase in ionisation energy from Z=2 to 10 and Z=11 to 18 relates to the increase in effective nuclear charge across periods 1 and 2 respectively.
  • The same valence shell is being filled across a period, hence inner shell shielding and nucleus – shell distance is approximately constant while the nuclear charge increases.
  • Hence the electron binding force increases, requiring greater energy to remove a valence shell electron. Therefore, the ionisation energy increases across a period.
  • When the next valence shell begins to fill, as occurs at Z=3, 11 and 19 , the ionisation energy drops substantially because the electron experiences increased inner shell nuclear shielding and is further from the nucleus.
  • The generally lower ionisation energies for the second period elements compared to the first period elements also reflects the filling of higher valence shells.
  • This reflects the decrease in ionisation energy down a group in the periodic table.

CHEMISTRY, M1 2013 HSC 20 MC

The structures of ozone and molecular oxygen are shown.
 

Ozone is more easily decomposed than molecular oxygen because

  1. it is polar.
  2. it is a bent molecule.
  3. it has a greater molecular mass.
  4. it has a lower average bond energy.
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`D`

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  • The presence of the single bond in the ozone molecule lowers the average bond energy.
  • Bond breaking is an endothermic process and so requires energy. As ozone has a lower average bond energy, less energy is required to break the bonds and thus decomposes more easily than molecular oxygen.

`=>D`