Christina SvenssonMr. Porter

SCH3UE- 03October 15th 2008

Trends in the periodic table

(a) The trend in ionization energy from top to bottom down a group

1) Ionization energy is the energy required to remove an electron completely from an atom.

I) Ionization energy

2. The elements present in 18th group correspond to the peaks of ionization energy; namely helium, neon and argon…

3. The alkalis (group 1) occupy the troughs on the graph; namely lithium, sodium, potassium….This observation is correlated to the electron configuration of these elements: they possess only one valence electron. As to obtain a configuration in octet (like the noble gases that precede them) they loose an electron easily to become a stable cation.

4. Within the same group of elements, the ionization energies decrease as the atomic number increases. 5. Across a row of elements, the ionization energy increases overall; however, when paying closer attention we notice that after the increase in ionization energy of two elements the third’s decreases. The pattern repeats itself until we the end of the period. II) atomic radii 2. The elements present in the alkali Group corresponds to the peaks in this graph; namely lithium, sodium, potassium….

3. The noble gas elements occupy the troughs of this graph. As we go across a period, electrons an additional electron is going into a sublevel of the same energy level, and protons are being added to the nucleus increasing the positive nuclear charge. As such, electrons are being pulled closer to the nucleus and the size of the atom decreases. Thus the radii of the atoms decrease as we move from the left hand side to the right and the noble gases have the smallest radii within their respective periods. 4. Within the same group of elements, the radius increases as the atomic number increases. 5. Across a row of elements, the radius decreases as the atomic number increases.

1) a) i) As we go from top to bottom of group IA, the nuclear charge increases.

ii) The ionization energy decreases as we go from top to bottom of a group.

b) i) As we go from top to bottom of group IA we notice that the atomic radius of the atom increases.

ii) This trend causes the ionization energy to decrease as we move from top to bottom of the group.

c) i) As we go from top to bottom of group IA, the shielding affect increases

ii) This trend causes the ionizing energy to decrease

Factors affecting ionization energy: top to bottom within a group

3. Down a group, the ionization energy decreases because the outermost electrons of the small atoms ( top of the group) are closer to the nucleus and more tightly held than the other electrons of a larger atom ( bottom of the group)

4. This line is slopping down this indicates an ionization energy decrease. This agrees with my previous conclusions.

1. a) i) As we go from left to right of period 2, the nuclear charge increases.

ii) This trend causes the ionization energy to increase

b) i) As we go from left to right across period 2, we notice that the atomic radius decreases

ii) This trend causes the ionization energy to increase.

increase/ decrease effect on I.E. (inc /dec)

nuclear charge↑ ↓

atomic radius↑ ↓

shielding effect ↑↓

c) i) As we go from left to right across period 2, the shielding affect stays constant

ii) This trend causes the ionization energy to decrease to remain relatively constant.

Factors affecting ionization energy from left to right across a period

3. Across a period, the ionization energy increases. The radii of atoms decrease as we move from the left hand side of the period to the right. Since the outermost electrons of the small atoms in a period are closer to the nucleus and are more tightly held than the outer electrons of a larger atom in a period the ionization energy increases as we move across a period.

4. This line is slopping up this indicates an ionization energy increase. This agrees with the conclusion of the above chart.

Additional Questions

1a) In period 4 Krypton (Kr) has the highest ionization energy b) In period 4 Potassium (K) has the lowest ionization energy

2) - Magnesium (Mg) has two valence electrons.- Aluminum (Al) has three valence electrons.

increase/ decrease effect on I.E. (inc /dec)

nuclear charge↑ ↑

atomic radius↓

↑

shielding effectRemains the same Remains relatively

constant

To achieve a complete octet of configuration like the noble gas that precedes them, the magnesium must loose 2 electrons and the aluminum must three electrons. As a result, the aluminum needs additional energy (the third ionization energy) to achieve an octet (as opposed to aluminum who only need up to the second ionization energy).

3. a) Between sodium and lithium, sodium is the most reactive. As we move down a group the radii of the elements becomes smaller. Ionization energy decreases as we move from the top to the bottom of a group because the outermost electrons of the small atoms of a group are closer to the nucleus and are more tightly held than the outer electron of a larger atom in the group. Hence, sodium has lower ionization energy than lithium meaning that he can loose his electron more readily. Therefore sodium is more reactive than lithium.

b) Between potassium and calcium, potassium is the most reactive. Alkali metals have lower ionization energy, in other words, they readily lose their single valence electron to achieve the octet, more so than calcium that need additional ionization energy. Because potassium loses his electrons more easily than calcium, he is the most reactive metal.

c) Between calcium and magnesium, calcium is the most reactive for the same reason as sodium (a)

d) Between strontium and rubidium, rubidium is the most reactive for the same reason as potassium (b)

Electron affinity1) Electron affinity is the energy given of when an electron is added to an atom.

A) Fluorine (F) is the element in period 2 with the highest electron affinity. As ionization energy increases within a period, electron affinity will increase in the same way. (neon is a noble gas it is not affected by electron affinity) Thus, fluorine has the highest electron affinity

B) Iodine (I) has the lowest electron affinity out these four elements. We have previously seen that the ionization energy decreases as we move down a group. As ionization energy decreases within a group, electron affinity will decrease in the same way. Thus iodine the lowest electron affinity.