Periodic Trends in Atomic Properties Why? Many properties of atoms have a repeating pattern when plotted with respect to atomic number. The similarities are due to the repeating pattern of electron configurations involving s, p, d, and f orbitals. The experimentally observed periodic trends in fact provide evidence for the orbital and shell structure of atoms and the meaningful arrangement of elements in the Periodic Table. Your ability to recognize the properties of the elements from their positions in the Periodic Table will prove useful in handling chemical compounds safely, developing new materials, finding new applications of known materials, or using chemistry in medical applications. Learning Objectives λ Develop relationships between position in the Periodic Table and electron configurations with atomic radius, ionization energy, and electron affinity. Success Criteria λ Use electron configurations to account for the relative sizes, ionization energies, and electron affinities of different atoms. λ Order elements by atomic radius, ionization energy, and electron affinity. Information Properties such as the size of an atom (atomic radius), the energy required to remove an electron from an atom (ionization energy), and the energy required to remove an electron from a negative atomic ion (electron affinity) can be understood in terms of the electron configuration of the atom and a competition between electron-nucleus attraction and electron-electron repulsion. An electron in an atom is attracted by the positively charged nucleus and repelled by the other electrons. Which one dominates depends on how effective the electrons are in getting close to each other or close to the nucleus. If the electron-nucleus attraction has a large effect, then the atom is small, has a high ionization energy, and large electron affinity. If the electron-electron repulsion has a large effect, then the atom is large, the ionization energy is small, and the electron affinity is very small, zero or negative. Resources Olmsted and Williams (Chemistry 3/e, Wiley, 2002) pp. 282-320. Prerequisites Coulomb’s law, atomic orbitals, Periodic Table, electron configurations New Concepts atomic radius, ionization energy, electron affinity, Coulomb attraction and repulsion Vocabulary first ionization energy, second ionization energy, periodic trend Periodic Trends in Atomic Properties
Definitions In your own words, write definitions of the terms in the New Concepts and Vocabulary sections.
Preliminary Activity
Complete the box diagrams below, as illustrated for carbon, to represent the atomic electron configurations.
1s 2s 2px 2py 2pz 1s 2s 2px 2py 2pz H C
He N
Li O
Be F
B Ne Periodic Trends in Atomic Properties
Model: Electron Ionization and Attachment and Electrostatic Interactions. Reaction Equations for Electron Ionization and Attachment. X → X+ + e− ∆E = 1st Ionization Energy X + e− → X− ∆E = − Electron Affinity
Dominant inter-particle forces in helium.
- e electron – nucleus attraction
electron – electron repulsion 2+
- electron – nucleus attraction e
Dominant inter-particle forces in lithium:
- e - e
3+
- e
Key Questions 1. What are the reaction equations associated with ionization energy and electron affinity?
2. From the model, why would you expect lithium, compared to helium, to be larger and have a higher ionization energy?
Periodic Trends in Atomic Properties
Model: Variation of Atomic Properties with Atomic Number The units are pm for atomic radii and kJ/mole for ionization energies and electron affinities.
250
200
150 Radii
ic 100 om
t 50 A 0 2500
2000 gies 1500
1000 on Ener i t
a 500
niz
o 0 I 350 300 es
i 250 200 finit f 150
on A 100 r t 50 ec l
E 0 0 5 10 15 20 25 30 35 40 45 50 55 Atomic Number Periodic Trends in Atomic Properties
Key Questions 3. As the atomic number increases across a row in the periodic table, does the electron – nucleus attraction increase or decrease?
4. Would the change in the electron – nucleus attraction that you identified in question 3 be expected to increase or decrease (a) the atomic radius?
(b) the ionization energy?
(c) the electron affinity?
5. As the number of electrons in an atom increases across a row in the periodic table, does the total electron-electron repulsion increase or decrease?
6. Would the change in electron-electron repulsion that you identified in question 5 be expected to increase or decrease the atomic radius, ionization energy, and electron affinity?
7. In going from hydrogen to helium, what do the changes in atomic radius (37 to 32 pm) and ionization energy (1311 kJ/mole to 2377 kJ/mole) suggest about the relative magnitudes of the changes in the electron – nucleus attraction and the electron – electron repulsion?
8. Given that the atomic radius and ionization energy of lithium are 152 pm and 520 kJ/mole, respectively, how does the orbital box diagram for lithium in the preliminary activity help to explain the observed differences in atomic radius and ionization energy between helium and lithium?
Periodic Trends in Atomic Properties
9. In moving across the second period of the Periodic Table, is the general trend in atomic radius consistent with the trend in going from hydrogen to helium?
10. For the ionization potentials in the second period, where do the exceptions to the general trend occur, and how do the orbital box diagrams in the preliminary activity help to explain these exceptions?
11. What happens to the atomic radii and ionization energies in going down a group, e.g., from Ne to Xe or Li to Rb?
12. How do the electron configurations of the atoms explain the trends you identified in question 11?
13. To what extent are the above trends in the atomic radius and ionization energy both across a row and between two rows found in other rows of the periodic table, and to what extent do new features occur?
14. Peaks in ionization energy occur at atomic numbers equal to 10, 18, 36, and 54, but these elements have very small or zero electron affinities. Peaks in the electron affinity occur at atomic numbers equal to 9, 17, 35, and 53. Why does a change of one unit for the atomic number result in such a drastic change in electron affinity? (Note: A high quality answer will provide and explanation that relates the electron configurations to the electron-electron repulsion and the electron-nucleus attraction.)
Periodic Trends in Atomic Properties
15. Why are the periodic trends in electron affinity and ionization energy more similar to each other than they are to the trend in atomic radii?
16. How can you use the periodic table and orbital diagrams to predict relative atomic radii, ionization potentials, and electron affinities of atoms?
Exercises 1. Explain why the first ionization energy of sulfur is less than the first ionization energy of phosphorous.
2. Explain why the first ionization energy of aluminum is less than the first ionization energy of magnesium.
3. Explain why is the electron affinity of Be very small or zero and the electron affinity of F very large (328 kJ/mole).
4. Identify the larger of each pair and explain why that one is larger. oxygen or sulfur
calcium or potassium
nitrogen or oxygen
copper or gold Periodic Trends in Atomic Properties
5. Identify which of the elements in each pair has the higher ionization energy and explain why for each. Ba or Cs
Br or Kr
S or Se
Si or C
6. Select the atom in each pair that has the greater electron affinity and explain why. S or Cl
C or O
Cl or Br
Si or P
S or Se
N or O