Atomic Electron Configurations and Chemical Periodicity

1. Atomic Electron Configurations and Chemical Periodicity Chapter 7

2. Dmitri Mendeleev • In 1870, Dmitri Mendeleev began to organize the periodic table due to repeating patterns - mass • one of the founders of the modern periodic table • 1913 Mosley arranged it by number of protons (atomic number)

3. 7.1 Pauli Exclusion Principle • No atomic orbital can contain more than two electrons and they must be opposite spin!

4. Pauli Exclusion Principle • Electrons can be represented by arrows in a box; this is referred to as an orbital box diagram. • Figure 7.2 describes the subshell filling order, or the order can be determined by using the periodic table.

5. 7.2 Atomic Subshell Energies • Electrons are assigned to subshells according to the aufbau principle. • assigned in order of “n + l” values • 2 subshells with same “n + l” value  electrons assigned to lower “n” first

6. Effective Nuclear Charge, Z* • Outer electrons may penetrate the inner electron region. These core electrons screen the positive nuclear charge from the outer (valence) electrons and so the outer electrons experience an average nuclear charge. • Electrons are found in configurations that result in the lowest energy for the atom.

7. 7.3 Electron Configurations • The use of orbital notations is referred to as the “spdf” notation. • The core electrons can be summed up in the “noble gas” notation. • What is the spdf for the lithium atom? • What is the orbital box diagram for the boron atom? • What is the noble gas notation for potassium?

8. Hund’s Rule • The most stable arrangement of electrons is that with the maximum number of unpaired electrons; minimizes electron-electron repulsions • All single electrons must have parallel spins to reduce repulsion.

9. Electron Configurations • Give the electron configuration of chlorine using the spdf, noble gas, and orbital box notations. • Write the electron configuration for Al using the noble gas notation and give a set of quantum numbers for each of the electrons with n = 3 (the valence electrons).

10. Electron Configurations • Transition Elements • Electrons may be found in s, p, and d sublevels • Most of the time, the periodic table can be used to determine electron filling • Differences may occur between the expected and the actual configurations… (not tested on the AP exam) • Chromium is expected to be [Ar]3d44s2; however due to the fact that the 3d and 4s are so similar in energy, each of the six valence electrons is assigned to a different orbital and therefore the actual configuration is [Ar]3d54s1 • Copper also has an unusual configuration of [Ar]3d104s1

11. Electron Configurations • Lathanides and Actinides • f subshells are filled or partially filled in the inner transition metals • La (lanthanum) [Xe]5d16s2 • Ce (cerium) [Xe]4f15d16s2

12. Practice Problems • What element is 1s22s22p63s23p5? • Do the spdf and orbital box diagram for phosphorus. • Do the spdf and noble gas for technetium (Tc) and osmium (Os).

13. Homework • After reading sections 7.1-7.3, you should be able to do the following… • p. 332 (11-21)

14. 7.4 Electron Configurations of Ions • Electrons are removed from the outermost energy level (shell of highest n). • If there are more than one subshell in the outermost level, the electrons are removed from maximum l • Na+ – the 3s1 electron is removed [1s22s22p6] • Fe2+ - [Ar]3d6

15. Practice Problem • What is the electron configuration of V2+, V3+, and Co3+? Are any of the ions paramagnetic? If so, how many unpaired electrons are there?

16. 7.5 Atomic Properties • The similarities in properties of the elements are the result of similar valence shell electron configurations. • Atomic Size • Atomic radius is ½ the experimentally determined distance between the centers of the two atoms. • For the main group elements, atomic radius increases going down a group due to the fact that the outermost electrons have a higher n value. • Atomic radius decreases going across a period due to the fact that effective nuclear charge increases as protons are added. • Transition metals are different due to the large filled d-sublevels; electron repulsion increases size toward the right.

18. Practice Problems • Place the three elements Al, C, and Si in order of increasing atomic radius. • If the interatomic distance in Br2 is 228 pm, what is the radius of Br? Using this value, and that for Cl (99 pm), estimate the distance between atoms in BrCl.

19. Ionization Energy • The energy required to remove an electron from an atom in the gas phase is referred to as ionization energy. • Excluding hydrogen, each atom can lose more than one electron and therefore has a series of ionization energies. • Each successive electron removal requires more energy because electrons are being removed from an increasingly positive ion.

20. Ionization Energy • For main group elements, first ionization energies increase across a period due to the increase in effective nuclear charge (increasing atomic number). • First ionization energies decrease down a group occurs because the electron removed is farther from the nucleus and therefore the nucleus-electron attraction is reduced.

21. Electron Affinity • The electron affinity, EA, of an atom is defined as the energy of a process in which an electron is acquired by the atom in the gas phase. • Both electron affinity and ionization energy represent the energy involved in the gain or loss of an electron by an atom. • An element with a high ionization energy generally has a high affinity for an electron.

22. Practice Problems • Compare the three elements B, Al, and C. • Place the three elements in order of increasing atomic radius. • Rank the elements in order of increasing ionization energy. • Which element is expected to have the most negative electron affinity value?

23. Ion Sizes • The radius of a cation is always smaller than that of the atom from which it is derived. Once an electron has been removed, the attractive force of the protons are exerted over fewer electrons. • Anions are always larger than the atoms from which they are derived due to the addition of electron(s) and an increase in electron-electron repulsions.

24. Ion Sizes

25. 7.6 Periodic Trends and Chemical Properties • Main group metals generally form cations with an electron configuration equivalent to that of the nearest noble gas. • Non-metals generally acquire enough electrons to form an anion with the electron configuration of the next, higher noble gas.

26. Homework • After reading sections 7.4-7.6, you should be able to do the following… • p. 333 (23-32)