Periodic Trends Atomic and Ionic Radius



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Periodic Trends
Atomic and Ionic Radius


  1. Explain why atomic size decreases from Na to Cl in the periodic table.

The trend, as you go across a period, is that the radius decreases. This is because all electrons are being added to the same shell (the same distance from the nucleus) and the nucleus is becoming increasingly positive due to addition of protons (effective nuclear charge increases). Since Cl has more protons in the nucleus, it pulls electrons towards it more tightly.


  1. Explain why the difference between the atomic radii of Na and K is relatively large compared to the difference between the atomic radii of Rb and Cs.

Sodium’s last electron is added to the 3rd energy level, K’s to the fourth. The difference between the size of Rb and Cs is between the fifth and sixth energy levels. The difference in energy of the lower shells is greater than the difference in energy of the higher shells. ;

Also, due to electron shielding the effect of the nucleus on the outer shell electrons is much greater for sodium than it is for Rb and Cs.




  1. Explain why a Ca atom is larger than a Zn atom.

Though Zn has more electrons than Ca, they were added to the third energy level, not the fourth. Both have the same number of outer electrons in the fourth energy level and Zn has a greater effective nuclear charge due to the larger number of protons in its nucleus. This makes Zn smaller than Ca.


  1. The radius of the Ca atom is 0.197 nm; the radius of the Ca2+ ion is 0.099 nm. Account for this difference.

Ca has two electrons in the fourth energy level (4s2), while the Ca2+ ion has lost these two electrons and has a full third energy level. Both have the same number of protons. Therefore, with fewer electron shells, the calcium ion will be smaller.


  1. Explain why the ionic radius of N3- is larger than that of O2-.

N3- has 7 protons and 10 electrons, O2- has 8 protons and 10 electrons. Both have identical electron configurations, but oxygen ion has a more positive nucleus and therefore pulls the electrons in a little more tightly.


  1. Ca2+ and Cl- are isoelectronic. Which has the larger radius? Explain why.

Ca2+ has 20 p and 18 e, Cl- has 17 p and 18 e. Electron configurations and shielding are the same, Ca2+ has a greater effective nuclear charge and therefore its radius will be a little smaller.
Ionization Energy




Ionization Energy (kJ / mol

First Second

K

419

3050

Ca

590

1140




  1. Explain the difference between Ca and K in regard to:

  1. Their first ionization energies

It takes more energy to remove the first electron from Ca than it does to remove one electron from K. This is because both electrons are being added to the same energy level (same distance from nucleus, same amount of shielding), but Ca has a more positive nucleus due to the additional proton so it holds its electron to it more tightly.


  1. Their second ionization energies

The second ionization energy of K is much higher than the second ionization energy of Ca. This is because Ca has two electrons in the 4th energy level while K has only one. The 2nd electron removed from K is removed from the third energy level. This energy level is closer to the nucleus so the electron is held much more tightly to the nucleus than the fourth energy level electron is. There is also more electron shielding between the nucleus and Ca’s 4s electron than there is between the nucleus and the 3rd energy level electron removed from K.


  1. The first ionization energy of Mg is 738 kJ/mol and that of Al is 578 kJ/mol. Account for the difference.

It takes more energy to remove the most energetic electron from Mg than it does to remove the most energetic electron from Al. It is slightly easier to remove the first p orbital electron than it is to remove the s orbital electron. This is because the p orbital is a little higher energy than the s orbital is those they are in the same n.


  1. Why is the first ionization energy of K lower than that of Li?

It takes less energy to remove an electron from K than it does from Li because the outer electon in K is in the fourth energy level and the outer electron of Li is in the second energy level. The K electron experiences a lower effective nuclear charge to the increased electron shielding and its greater distance from the nucleus.


  1. Below are ionization energies for third period elements.


First Ionization Energy (kJ/mol)

Second Ionization Energy (kJ/mol)

Third Ionization Energy (kJ/mol)

Element 1

1251

2300

3820

Element 2

496

4560

6910

Element 3

738

1450

7730

Element 4

1000

2250

3360




  1. Which element is most metallic in character?

The most metallic element is located furthest to the bottom left of the periodic table. Since all elements are in the same period, the most metallic is the one closest to the left. Element 2 has the lowest first ionization energy so it must have the fewest protons in its nucleus and is further to the left of the period. Also it has a very large increase between its first and second ionization energy which means the second electron is removed from a lower energy level and it has only one electron in its outer shell.


  1. Identify element 3. Explain your reasoning.

Element 3 must be Mg. The very large increase in ionization energies is between removal of the second and third electrons. This means the element has two electrons in its outer energy level and the third electron is removed from the second energy level. Mg has 3s2 electrons in the third energy level.


  1. Write the electron configuration for element 3.

1s22s22p63s2


  1. What would be the expected valence number for the most common ion of element 2?

Element 2 has one electron in the third energy level (see part a) and it will lose this electron to form an Na+1 ion which is isoelectronic with Ne.


  1. What is the chemical symbol for element 2?

Na


  1. A neutral atom of which of the four elements has the smallest radius?

The element to the right of the period will have the most protons in its nucleus and this greater effective nuclear charge will cause it to pull its electrons in a little more closely and have the smallest size. Since element 1 has the highest first ionization energy, it has the most positive nucleus.
Electron Affinity
11. Which element has the most negative electron affinity: B, Al, C, Si? Explain why.

The element with the most negative electron affinity releases the most energy when an electron is added to the neutral atom and therefore has the greatest electron affinity. The element with the greatest attraction for an electron, is pulling an electron in at a lower energy level, closer to the nucleus where there is less electron shielding. It will also have more protons in the nucleus. Since Al and Si are adding an electron to energy level 3, they will have lower electron affinity than B and C (energy level 2). Carbon has more protons in its nucleus than B and this greater effective nuclear charge will give it a higher electron affinity.


12. Why are the electron affinities of group 2A elements lower than 1A elements (why don’t they follow the expected trend)?

Adding an electron to group 2A elements would be placing the first electron in the p orbital. Since p orbital electrons are slightly higher energy, the electron affinity for these electrons is a little lower than it is for the s electron.


13. Why are the electron affinities of group 5A elements lower than 4A elements (why don’t they follow the expected trend)?

Adding an electron to group 5A elements would be adding the first paired p electron. Because there is a slight repulsion between like charged electrons, the electron affinity for this electron is a little lower than it is for the unpaired p electrons.


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