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By referring to their electron configurations, explain why metals in the first transition series of the periodic table typically have several different oxidation states while Group II metals have only one - VCE - SSCE Chemistry - Question 6 - 2004 - Paper 1
Question 6
By referring to their electron configurations, explain why metals in the first transition series of the periodic table typically have several different oxidation sta... show full transcript
Worked Solution & Example Answer:By referring to their electron configurations, explain why metals in the first transition series of the periodic table typically have several different oxidation states while Group II metals have only one - VCE - SSCE Chemistry - Question 6 - 2004 - Paper 1
Step 1
a. By referring to their electron configurations, explain why metals in the first transition series of the periodic table typically have several different oxidation states while Group II metals have only one.
Answer
The first transition series metals have complex electron configurations, with partially filled 3d and 4s subshells. Due to the closeness in energy levels of these subshells, electrons can be lost from both the 4s and 3d orbitals, leading to the ability to achieve multiple oxidation states. For example, iron (Fe) can exhibit +2 and +3 oxidation states by losing different combinations of electrons from these subshells.
On the other hand, Group II metals, such as magnesium and calcium, have a straightforward electron configuration with a filled s subshell (ns²). They typically lose the two outermost s electrons to form a +2 oxidation state. This lack of additional available electrons from a d subshell limits their oxidation states to only one.
Step 2
b. Predict the trend in the following properties of the second period elements, moving from lithium to fluorine, giving an explanation in support of each prediction.
Answer
| Property | Increases or decreases? | Explanation for predicted trend |
|---|---|---|
| Atomic size | Decreases | As we move across the period, the nuclear charge increases, drawing the electrons closer to the nucleus and reducing atomic size. |
| Electronegativity | Increases | With an increase in nuclear charge, the attraction of the nucleus for the outer shell electrons strengthens, leading to higher electronegativity. |