Supplemental Activities

Module: Atomic Theory

Section: Periodic Properties and Trends - Key

Periodic Table and Reactivity

Activity 1

1. Consider lithium metal. a. Why don’t we find lithium metal in its neutral atomic state in nature?

The one valence electron that lithium metal has is easily given up so that lithium metal can return to a noble gas state. Neutral lithium metal is highly reactive because it’s energetically favorable for it to give up its one valence electron.

b. In what electron configuration would you expect to find lithium metal?

We would expect to find lithium with one less electron. So instead of its neutral electron configuration (1s22s1), we would expect to see it in the configuration 1s2.

c. Comment of the similarities between your answer for part b and the electron configuration for helium.

Helium has the electron configuration 1s2. This is the same electron configuration that we would expect to find for lithium in nature. They would both have two valence electrons that fill up the n = 1 energy level.

2. Why do groups/families tend to share common chemical behaviors? In your explanation, discuss electron configurations and valence electrons.

We can observe similar patterns in the electron configurations of elements in a group. As we can see in these electron configurations, the elements down a group share the same number of valence electrons. Therefore, their bonding habits and chemical reactivity will be similar.

Activity 2

1. There is a radioactive isotope of strontium that can cause bone and blood cancers. Why would strontium replace calcium in human bones (discuss valence electrons)?

Strontium and calcium are in the same family of elements (alkaline earth metals). They have the same number of valence electrons (two) so they interact in nature in the same way.

2. Find a noble gas and a common ion that are both isoelectronic with Sr2+and write down the electron configuration that these species all share.

The noble gas that is isoelectronic with Sr2+ is Kr. There are a few choices for common anions that are isoelectronic with Sr2+ : Br–, Se2–, As3– The electron configuration for all these species is: 1s22s22p63s23p64s23d104p6

Activity 3

1. Write down the charge the different groups tend to make when forming ions and explain why: a. Group IA – +1 b. Group 2A – +2 c. Group 5A – –3 d. Group 6A – –2 e. Group 7A – –1 f. Why? Elements in these groups will either lose or gain electrons and therefore attain positive and negative charges respectively. These cations (positive) and anions (negative) are more stable than their neutral states because these loss or gain of electrons allows them to achieve a noble gas electron configuration state.

2. How is the effective nuclear charge of an atom different than its nuclear charge? Why is there a difference?

Nuclear charge is simply the charge of the nucleus of an atom. It is equivalent to the number of protons in the nucleus. So carbon is +14, nitrogen is +15 and so on. Effective nuclear charge, however, is the charge the valence electrons feel from the nucleus. The core electrons (which are negatively charged) are shielding the valence electrons from the positive nucleus. The valence electrons of a neutral atom (equal number of protons and electrons) will feel a decreased nuclear charge due to shielding and this is the effective nuclear charge. An exception is a neutral atom of hydrogen because it has no core electrons and so the one valence electron feels no repulsive interference for other electrons within the atom.

3. How does effective nuclear charge change as you move down a group of elements? As you move left to right across a period of elements?

In order to calculate the effective nuclear charge of an element, subtract the number of core electrons from the nuclear charge (number of protons). As you move down a group of elements the value of the effective nuclear charge stays the same. However, the valence electrons are further and further away from the nucleus and this is something to consider when discussing periodic trends. As you move left to right, the effective nuclear charge increases. Electrons are being added to the same energy level as the valence electrons and therefore do not contribute to the shielding effect. At the same time the number of protons in the nucleus is increasing. So the effective nuclear charge increases from left to right across a row. 4. Consider the element calcium (Ca). What is the effective nuclear charge experience by electrons in the following of Ca’s subshells? a. 4s 2 b. 3p 10 c. 3s 10 d. 1s 20

5. If a helium atom had an electron had one of its electrons excited from its 1s subshell to a 2s sbshell, what would be the ENC experienced by that electron?

He has two protons, and the unexcited electron would now be screening the excited electron, so it would experience an ENC of: 2 - 1 = 1.

6. Consider an electron in the orbitals given and rank them from most attracted to the nucleus to least attracted to the nucleus by comparing their effective nuclear charges that they each experience. (i) 3p in Mg (ii) 2p in O (iii) 4s in Ca (iv) 2s in K iv > ii > i> iii

Periodic Trends

Activity 1

1. As you move down a group what do we find about the ionization energy, electron affinity and atomic radii? How about as you move across a period from left to right? From top to bottom: ionization energy and electron affinity decrease whereas the atomic radii increase. From left to right: ionization energy and electron affinity increase whereas the atomic radii decreases. Arrange the following groups of atoms in order of decreasing size:

2. P, Sb, N, As – Sb > As > P > N a. Br, V, N, Ba – Ba > V > Br > N Whicb.h atom or ion in each of the following lists has the largest ionization energy: 3. Mg, S Ba – Mg (they are all in the same group, Mg is the highest up, is the smallest and has the highest ionization energy) a. r, Ca, Co, Se – Se (they are all in the same period, Se is the farthest to the right, is the smallest and therefore has the highest ionization energy) b. O2–, O, O2+ – O2+ (they all have the same number of protons, but O2+ has the fewest number of electrons and is the smallest. Therefore the effective c. nuclear charge is greater and it requires more energy to remove an electron)

Activity 2

1. Explain how shielding affects ionization energy and atomic radii. The shielding of outer electrons by the core electrons decreases the ability of the nucleus to pull on the outer electrons. When valence electrons are added to an entirely new shell as you move down a group, the effective nuclear charge (Zeff) stays the same but the outer electrons are far removed from the nucleus. Therefore the atomic radii increase and ionization energy decreases as we move down a group. However, as we move across a period from left to right, electrons are added to the same valence shell. The shielding isn’t as effective because the same number of core electrons is shielding the valence electrons from an increasingly larger nuclear charge. The effective nuclear charge on the valence electrons increases and therefore the atomic radii decrease and ionization energy increases as we move across a period from left to right. 2. Here is a way to depict the process related to the definition of electron affinity using the element nitrogen: N (g) + 1e–  N– (g) How could you depict the process related to the definition of ionization energy in a similar way using the element nitrogen? Ionization Energy: N (g)  N+ (g) + 1e– The definitions of both electron affinity and ionization energy specify that the addition or removal of an electron when the element is in the gas phase. So the value for ionization energy of a particular element depends on the phase that the element is in when the electron is removed. 3. Why do noble gases have exceptionally low electron affinities and exceptionally high ionization energies?

Because noble gases have completely filled shells, which are highly stable, they have little energetic motivation to acquire or release electrons. They’re rather quite happy just as they are.

4. Write the electron configuration for silver. Try to give at least one example of an ion that would have an identical electron configuration.

[Kr] 5s1 4d10 Cd+, In2+ and Sn3+ would all have the same electron configuration as Ag.