Trends In The Periodic Table. C.s.

Christina Svensson Mr. Porter SCH3UE- 03 October 13th 2008

Trends in the Periodic Table

1

Part I) Ionization energies and atomic radius: 1) Define ionization energy for an element. The energy required to remove the outer electron from a gaseous atom (all by itself; not connected to other atoms in a solid or liquid). Y= Ionization energy, X= Atomic radii. See graphs attached. 2) What groups of elements corresponds to the peaks of y/x on the graph? x) The elements in the 18th group; Helium, Neon, Argon, Krypton, Xenon, Radon. y) The elements in the 1st group; Lithium, Sodium, Potassium, Rubidium, Cesium and Francium. 3)

What group of elements occupies the troughs on the graph of atomic radius? Suggest a reason for this. x) The alkalis, which you find in group one; Lithium, Sodium, Potassium, Rubidium, Cesium, Francium. A reason for this would be that the only have one valence electron, and to obtain a full outer shell (eight electrons), they loose an electron. y) The noble gasses occupy the troughs on the graph. Because an additional electron is going in to a sublevel of the same energy level, and causes the protons to be pulled closer to the nucleus and increasing the positive nuclear size. And by pulling the protons closer, the electrons will also be pulled closer to the nucleus, and the atom radii decreases.

4)

What trend is seen in atomic radii as the atomic number increases across a group of elements? x) Within the same group of elements, the ionization energies decrease as the atomic number increases. y) Elements within the same group increase its radii as the atomic number increases.

5)

What trend is seen in atomic radii as the atomic number increases across a period or row of elements? x) In general, the ionization energy increases, but if we take a closer look we see that there is a pattern: two elements ionization energy increases, and that is the third ones decrease. This pattern is repetitive trough the period. y) Elements within the same row is decreasing its radii and increasing its atomic number.

2

Part II) The trend in ionization energy from top to bottom of a group: 1) As you go from top to bottom down group 18 (use this as your e. g.): a. I) Does the nuclear charge increase or decrease? It increases. II) Does this trend cause the ionization energy to increase or decrease? It decreases. b.

I) Does the atomic radius increase or decrease? It increases. II) Does this trend cause the ionization energy to increase or decrease? It decreases.

c.

I) Does the shielding effect increase or decrease? It increases. II) Does this trend cause the ionization energy to increase or decrease? It decreases.

2) increase/ decrease

effect on I.E. (inc /dec)

↑

↓

↑

↓

nuclear charge atomic radius shielding effect

↑ ↓

3) Going down a group, the ionization energy decreases because the electrons of the smallest atoms the furthest away are closer to the nucleus and more tightly held in place than other electrons from a larger atom (bottom of group) 4) It indicates an increase in ionization energy. Yes, it agrees with my previous conclusions. B1) aI) Increases. II) Increases. bI) Decreases. II) Increases. cI) Stays constant. II) Stays relatively constant.

B2) 3

increase/ decrease

effect on I.E. (inc /dec)

↑

↑ ↑

nuclear charge atomic radius ↓ shielding effect Remains the same

Remains relatively constant

3) The ionization energy increases across a period. The radii of the atoms decrease as we go from left to right. And because the outer electrons of the small atoms in a period are closer to the nucleus, and are more tightly attached than the outer electrons in a larger atom, the IE increases as we go across the period. 4) This indicates an I. E. increase. And yes, it agrees with the conclusion drawn from the upper chart.

ADDITIONAL QUESTIONS:

4

1a) In period 4 Krypton has the highest I. E. b) In period 4 Potassium has the lowest I. E. 2)

 Magnesium has two valence electrons.  Aluminium has three valence electrons. To achieve the octet (as the noble gasses) making the atoms stabile, Mg must loose two electrons and Al three. Because of this, Al needs additional energy (the 3rd I. E.) and Mg needs the 2nd I. E. to achieve the octet. See drawing attached. 3a) Between sodium and lithium, sodium would be the most reactive. When moving down a group, the radius of the elements is getting smaller. At the same time, I. E. decreases as we move from top to bottom of a group, as explained in previous questions; 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. And therefore sodium has a lower I. E. than lithium, which means that Li can loose its electrons easier than Na. Therefore Na (Sodium) is the most reactive. b) Between potassium and calcium, potassium is the most reactive. This is because Alkali metals have a lower I. E., meaning that they easier loose their electrons to form an octet, so calcium needs additional I. E.. And because Ka looses its electrons easier, it is the most reactive metal. c) Between calcium and magnesium, calcium is the most reactive. See #a for explanation. d) Between strontium and rubidium, rubidium is the most reactive. See #b for explanation. Electron affinity: 1) Electron affinity is the energy given when an electron is added to an atom. A) In period two, fluorine is the element with the highest electron affinity. It is the same as the I. E.; as it increases, electron affinity will increase in the same way. B) Iodine has the lowest electron affinity out of these elements. As the I. E. decreases as we move down a group, the electron affinity will decrease in the same way. 2)Drawings

Graphs

5

radius versus atomic number 250

200

radii

150

100

50

0 0

5

10

15

20

25

atom ic num ber

Ionization Energy vs Atomic Number 2400

Ionization Energy Kj/Mole

2000

1600

1200

800

400

0 0

5

10

15

20

25

Atomic Number

6