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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 13

Using your knowledge of the atomic radii for the 2nd and 3rd period elements.

  1. Explain the general trend in atomic radius across Period 2.   (3 marks)
  1. Explain why the elements of Period 3 have larger atomic radii than the corresponding elements of Period 2.   (2 marks)
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a.   Across Period 2, atomic radius decreases from lithium to neon:

  • This occurs because the number of protons (nuclear charge) increases across the period, while electrons are added to the same energy level.
  • As a result, the increased nuclear attraction pulls the electron cloud closer to the nucleus.
  • This leads to reduced atomic radius despite more electrons being present.
  • Therefore, there is a clear trend of decreasing atomic size from left to right across Period 2.

b.   Period 3 atomic radii are larger than corresponding Period 2 element:

  • Period 3 elements have larger atomic radii than Period 2 elements because they have one additional electron shell.
  • This occurs because Period 3 elements have electrons in the n=3 shell, while Period 2 elements only fill up to n=2.
  • Period 3 elements have more protons that create a greater nuclear attractive force with the orbiting electrons but the extra electron shell outweighs the effect of increased nuclear charge.
  • As a result, the outermost electrons in Period 3 are further from the nucleus, causing increased atomic size despite having more protons.
Show Worked Solution

a.   Across Period 2, atomic radius decreases from lithium to neon:

  • This occurs because the number of protons (nuclear charge) increases across the period, while electrons are added to the same energy level.
  • As a result, the increased nuclear attraction pulls the electron cloud closer to the nucleus.
  • This leads to reduced atomic radius despite more electrons being present.
  • Therefore, there is a clear trend of decreasing atomic size from left to right across Period 2.

b.   Period 3 atomic radii are larger than corresponding Period 2 element:

  • Period 3 elements have larger atomic radii than Period 2 elements because they have one additional electron shell.
  • This occurs because Period 3 elements have electrons in the n=3 shell, while Period 2 elements only fill up to n=2.
  • Period 3 elements have more protons that create a greater nuclear attractive force with the orbiting electrons but the extra electron shell outweighs the effect of increased nuclear charge.
  • As a result, the outermost electrons in Period 3 are further from the nucleus, causing increased atomic size despite having more protons.

CHEMISTRY, M1 EQ-Bank 11

Explain the relationship between electronegativity and atomic radius with non-metal reactivity down Group 17 (the halogens) of the Periodic Table.   (4 marks)

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  • Atomic radius increases down Group 17 because each element gains more electron shells.
  • This leads to weaker nuclear attraction for incoming electrons, due to greater distance from the nucleus and increased shielding by inner shells.
  • As a result, electronegativity decreases down the group, and consequently decreasing reactivity down Group 17
  • For instance, fluorine (small radius, high electronegativity) is most reactive while in contrast, iodine (large radius, low electronegativity) is least reactive.
  • Therefore, smaller atoms with higher electronegativity are more reactive non-metals.
Show Worked Solution
  • Atomic radius increases down Group 17 because each element gains more electron shells.
  • This leads to weaker nuclear attraction for incoming electrons, due to greater distance from the nucleus and increased shielding by inner shells.
  • As a result, electronegativity decreases down the group, and consequently decreasing reactivity down Group 17
  • For instance, fluorine (small radius, high electronegativity) is most reactive while in contrast, iodine (large radius, low electronegativity) is least reactive.
  • Therefore, smaller atoms with higher electronegativity are more reactive non-metals.

CHEMISTRY, M1 EQ-Bank 10

Using your knowledge of electronic configurations, explain what happens to atomic radii as you go across a period from left to right in the periodic table.

Identify one element from the first period with a larger atomic radius and one with a smaller atomic radius than Boron.   (3 marks)

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  • As elements progress from left to right across a period, each successive element has one more proton and electron than the previous one.
  • These additional electrons enter the same energy level, increasing the effective nuclear charge experienced by electrons.
  • However, the stronger pull from the nucleus due to the higher positive charge causes the electrons to be drawn closer, resulting in a decrease in atomic radius.
  • Lithium has a larger atomic radius than Boron, while Neon has a smaller atomic radius than Boron.
Show Worked Solution
  • As elements progress from left to right across a period, each successive element has one more proton and electron than the previous one.
  • These additional electrons enter the same energy level, increasing the effective nuclear charge experienced by electrons.
  • However, the stronger pull from the nucleus due to the higher positive charge causes the electrons to be drawn closer, resulting in a decrease in atomic radius.
  • Lithium has a larger atomic radius than Boron, while Neon has a smaller atomic radius than Boron.

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 EQ-Bank 5

"Electronegativity increases as you move across periods left to right and decreases as you move down groups".

Explain this trend with reference to the following periodic table.   (4 marks)
  

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Moving from left to right:

  • As you move across a period from left to right in the same row, the number of protons in the nucleus of elements increases in accordance with their atomic number.
  • eg. \(\ce{Li}\) (far left) has 3 protons in its nucleus whereas \(\ce{F}\) (far right) has 9 protons in its nucleus.
  • This leads to a greater attractive force and thus higher electronegativity. 

Moving down within a group (column):

  • Adding electron shells to a nucleus decreases electronegativity.
  • This is due to an increase in atomic radius and the effect that extra electron shells have in shielding the attractive forces of protons.
  • All the elements in a period (row) further down the periodic table have an extra electron shell than the period directly above them, decreasing electronegativity as you move down.
Show Worked Solution

Moving from left to right:

  • As you move across a period from left to right in the same row, the number of protons in the nucleus of elements increases in accordance with their atomic number.
  • eg. \(\ce{Li}\) (far left) has 3 protons in its nucleus whereas \(\ce{F}\) (far right) has 9 protons in its nucleus.
  • This leads to a greater attractive force and thus higher electronegativity. 

Moving down within a group (column):

  • Adding electron shells to a nucleus decreases electronegativity.
  • This is due to an increase in atomic radius and the effect that extra electron shells have in shielding the attractive forces of protons.
  • All the elements in a period (row) further down the periodic table have an extra electron shell than the period directly above them, decreasing electronegativity as you move down.

CHEMISTRY, M1 EQ-Bank 4

Describe two factors which affect the degree of electronegativity of an atom.   (2 marks)

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Number of Protons in the Nucleus:

  • The greater the number of protons in the nucleus of an atom, the greater the attractive force exerted on electrons and therefore the greater the electronegativity.

Atomic Radius:

  • The smaller the atomic radius, the larger the attractive force that can be exerted on electrons (due to lack of shielding).
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Number of Protons in the Nucleus:

  • The greater the number of protons in the nucleus of an atom, the greater the attractive force exerted on electrons and therefore the greater the electronegativity.

Atomic Radius:

  • The smaller the atomic radius, the larger the attractive force that can be exerted on electrons (due to lack of shielding).

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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  • 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`