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CHEMISTRY, M1 EQ-Bank 31

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.

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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.

CHEMISTRY, M1 EQ-Bank 28

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

    • atomic radii
    • first ionisation energy
    • electronegativity

Justify 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 14

"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 13

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`