This question seems too simple to me but i was just wondering if someone could help me with it?
8 a) Explain why the first ionisation energy values for the Group 1 elements become less positive as you go down the group? I have done this part it's b which is what I would like checked please.
b) Why is this called a trend?
Is it: This is called a trend because as you go down the group the positivity also goes down (devreases).
I looked up the word "trend" in the dictionary and came up with the following:
a general direction in which something is developing or changing."
If I might offer a suggestion to make the words flow better---
This is called a tend because it becomes easier to ionize the atom as one goes down the periodic table within a group.
*decreases
and ok, thank you!
In terms of the explanation for part b), it is not entirely accurate. While it is true that the first ionization energy values for Group 1 elements become less positive as you go down the group, the term "trend" refers to the overall pattern or trend that is observed when examining a group of elements or a series of data points.
In the context of ionization energy values, the term "trend" specifically refers to the general tendency observed when comparing the values of the first ionization energy across the elements in a group. In this case, the trend is that the first ionization energy generally decreases as you go down the Group 1 elements.
The trend in ionization energy is a result of two main factors:
1. Increasing atomic radius: As you go down the Group 1 elements, the atomic radius increases due to the addition of more energy levels (electron shells). The increased distance between the outermost electron and the nucleus weakens the attractive force, making it easier to remove the outermost electron. Therefore, the ionization energy decreases.
2. Shielding effect: As you go down the Group 1 elements, the number of inner-shell electrons also increases. These inner-shell electrons act as a shield, reducing the effective nuclear charge experienced by the outermost electron. Consequently, the electron is less strongly attracted to the nucleus, leading to a decrease in ionization energy.
So, to summarize, the decreasing trend in the first ionization energy values for Group 1 elements as you go down the group is due to the combined effects of increasing atomic radius and the shielding effect of inner-shell electrons.