Ionization energy, Atomic Radii and Electronegativity Periodic Trends Ionization energy, Atomic Radii and Electronegativity

What Holds Electrons in orbitals around the Nucleus? Their (electrostatic) attraction for the positive PROTONS in the nucleus. Their spin holds them out from the nucleus and stable in their orbits.

1. Ionisation energy Definition: The energy required to remove the least tightly bound electron from each atom of one mole of gaseous atoms. Measured in kJ mol-1 Second ionization energy greater than the first as it is much harder to remove electrons from a positive ion.

Is ionisation energy Endo or Exo thermic? Always ENDOTHERMIC as energy is required to remove electrons from an atom or ion.

Successive Ionisation energies. Look at your handout showing ionization energies for potassium as more electrons are removed. Why is the graph not smooth? After the first electron is removed from the fourth shell (outside shell), we come to the third shell which is closer to the positive nucleus so it is harder to remove electrons, so more energy is required. After the 9th electron is removed, we come to second shell which is even closer to the nucleus and even more difficult.

Ionisation energies for the first twenty elements – what do you notice?

First ionization energies of the elements. In general, as we move from left to right across the P.T., ionization energies ______________. Elements same number of shells but nuclear charge increase ( No protons in the nucleus) And as we move down the periodic table, ionisation energies ____________. Elements have more shells, valence electrons are further away less energy required However, when we look at a graph of first ionization energies, we can see that some elements are more stable than others. Be is more stable than B and N is more stable than O. (Write the electron configurations of these elements then explain why some are more stable.)

First ionization energies of the elements. In general, as we move from left to right across the P.T., ionization energies increase. Elements have same number of shells but nuclear charge increases ( More protons in the nucleus) And as we move down the periodic table, ionisation energies decrease Elements have more shells, valence electrons are further away less energy required However, when we look at a graph of first ionization energies, we can see that some elements are more stable than others. Be is more stable than B and N is more stable than O. (Write the electron configurations of these elements then explain why some are more stable.)

Answers Be – 1s22s2 B- 1s22s22p1 N – 1s22s22p3 O – 1s22s22p4 Be is more stable than B as it only has a full 2s shell whereas B has a single electron in its 2p shell which is less stable and more easily removed. Thus it has a lower ionisation energy than Be N is more stable than O as it has one electron in each of it’s 2p orbitals which is more stable than 4 electrons in the 2p level. Thus N has a higher ionisation energy than O and is more stable.

Note: Greater the ionisation energy the more difficult it is to remove the electron. Additional shells means greater shielding effect between valance shell and nucleus. So less energy required to remove outer most electron.

2. Atomic radii How BIG the atom or ion is It is impossible to measure an atomic radius of an individual atom. It is possible to measure the distances between adjacent nuclei of various substances. The radii of noble gases can’t be measured - they don’t form bonds.

TRENDS IN ATOMIC RADIUS DOWN A GROUP Going down a group each atom has the same numbers of valence electrons and the same effective nuclear charge…. But there are more shells as we move down the group so the valance shell is further from the nucleus So effective nuclear charge felt is less Valence layer is in effect shielded from the nucleus by inner electrons so the electrostatic force between the nucleus and valence electron decreases

Trends in ionic size/radii across a row As we go across the PT the number of electrons in the shell increases and so does the number of protons o do the numbers of protons in the nucleus. The electrons are all the same distance from the nucleus so each is attracted to the positive nucleus by the same amount. However, as the numbers of protons increases, the electrostatic attraction ( effective nuclear charge) for each electron increases so that the shell is pulled in closer.

ATOMIC RADIUS summary Radii ________ going down a group due to increasing numbers of _________ __________. Radii __________ going across a period however. Why are the electron shells pulled closer as more electrons are added?

IONIC radii – cations - blue circles

The trend… Cations are ___________ in size than their respective atoms. This is because to become cations they lose electrons so they have a stable outer shell. In doing so, they lose a whole shell so the size is smaller. AND…The nuclear charge is overall very positive for the numbers of electrons so all electrons are drawn into the nucleus tightly.

Ionic radii – anions – green circles

Anion trend Anions are __________ than their respective atoms. To become an anion they must _________ electrons. In doing so they gain a full outer shell. There are more electrons in the outside shell than before and as electrons _______ each other they need more room and spread out so that the radius is bigger. AND… there is reduced attraction for the __________ nucleus as the overall positive charge is less than the ___________ charge of the electrons so they are not held so tightly to the nucleus.

Question 1. Compare the relative sizes of the F atom and F- ion and explain the difference in their radii.

Answer The F atom has a smaller radii than the F- atom. F has an electron configuration of 1s22s22p5 and a nuclear charge of + 9, F- has a configuration of 1s22s22p6 and a nuclear charge of +9. The 7 electrons in the outer shell of F have a strong attraction to the highly positive nucleus so are pulled in close. However, there are 8 electrons in the outer shell of F-. These have a reduced attraction for the nucleus as the positive charge of the nucleus must now attract 8 electrons rather than 7 and the extra electron causes them to repel each other and spread out more. Thus F- is larger than F atom. .

Question 2 Compare the relative sizes of the F- and Na+ and explain the difference in their radii.

Answer F- and Na+ have the same electron configuration, 1s22s22p6. However, F­ – is much larger than Na+ as it has a much weaker nuclear charge(+9) than Na+ (+11) so it does not hold it’s outer shell electrons as tightly or close and this means the radius of the F- ion is larger.

ELECTRONEGATIVITY The measure of the attraction an atom has for bonding electrons Fluorine has the highest followed by oxygen and nitrogen

Electronegativity _________ as you go across a period charge and decreases as you go down a group It is measured by a combination of ionisation energy, electron affinity and bond dissociation energies. It increases across a period as the increasing ___________ charge attracts electrons more. It decreases down a period as electron _______are added and this shields the effect of the nuclear ___________, reducing the attraction for electrons.

Electronegativity Increases as you go across a period charge and decreases as you go down a group It is measured by a combination of ionisation energy, electron affinity and bond dissociation energies. It increases across a period with as the increasing nuclear charge attracts electrons more. It decreases down a period electron shells are added and this shields the effect of the nuclear charge, reducing the attraction for electrons.