Chapter 8: Electron Configuration and Chemical Periodicity

1. Chapter 8: Electron Configuration and Chemical Periodicity

2. Periodicity

3. Development of the Periodic Table

4. Electron Spin In 1928, it was discovered that an electron has an intrinsic angular momentum, or spin. In a magnetic field, the rotation axis has only two possible orientations.

6. Multi-electron Atoms

7. Effective Nuclear Charge We account for electron repulsion by assuming that the electrons shield each other from the nuclear charge. The net nuclear charge experienced by an electron is the effective nuclear charge, Zeff. If S is the average number of screening electrons: Zeff = Z - S

9. Electron Energies Due to screening, different subshells have different energies, increasing in the order: s < p < d < f

10. Plan for the Day Warm-up to the Atom again! The Periodic Table and Atomic structure Models of Chemical Bonding Nomenclature

11. The Exclusion Principle How many electrons can fit in, or “occupy” an orbital? The Pauli Exclusion principle states: No two electrons in an atom can have the same four quantum numbers. The ground state of helium has two electrons in the 1s orbital, but with opposite spins. n l ml ms electron 1 1 0 0 +½ electron 2 1 0 0 -½

12. Many-Electron Atoms The Aufbau principle: Electrons are assigned, one at a time, to hydrogen orbitals with lowest possible energy. An orbital diagram shows the number of electrons in each occupied orbital.

13. Hund’s Rule Hund’s rule: The lowest energy state has the most unpaired electrons. Carbon:

14. Energy diagrams for each atom in first two rows

15. Locating Electrons with Quantum Numbers 1. Draw the orbital energy diagram for the Boron atom. 2. Write a set of quantum numbers (4) for each of the electrons of B.

16. Conceptual Question Identify the subshell in which electrons with the quantum numbers n = 6, l = 1 may be found. a)   5pb)   6dc)   6pd)   6fe)   3d

17. Conceptual Question What is the lowest-numbered principal shell in which f orbitals are found? 2 1 4 5 3

18. Concept Question: Which of the following sets of quantum numbers are allowed for an electron in an atom? n l ml ms1) 2 1 0 +1/22) 3 0 +1 -1/23) 3 2 -2 -1/24) 1 1 0 +1/25) 2 1 0 0 a)   2, 4b)   1, 3c)   3, 4d)   1, 2, 3e)   2, 4, 5

19. Conceptual Question Which of the following electron diagrams represents a correct ground state? 

20. Electron Configurations and the Periodic Table The electron configuration of an atom can be estimated from the Periodic table. The actual configuration must be determined by experiment.

21. Electron Configurations and the Periodic Table Write electron configurations for:

23. Elements within a group have similar electron configurations in their valence electron shells Alkali metals 1s1 2s1 3s1 4s1

24. Nobel Gases Filled valence shell – HAPPY!

25. Valence Shell/Electrons Valence shell: outer most energy level/shell that contains electrons in an atom Valence Electrons: electrons in the outermost shell Core electrons: Inner electrons (non-valence electrons) Examples C Rb Br

26. Activity Create a model of an O atom. Use all that you’ve learned about modern atomic structure including the location of the subatomic particles, orbital drawings (3D), orbital energy diagram, electron configuration, quantum numbers, valence electrons, core electrons, etc… What other atoms in the periodic table have a similar electron configuration to O?

27. Periodic Properties of the Elements Early versions of the Periodic table were constructed by Mendeleev and Meyer. We now know that the periodic properties are due to the electronic structure of atoms. Electronic structure explains the observed trends in Atomic size Ionization Energy Electron Affinity Å

28. Effective Nuclear Charge We account for electron repulsion by assuming that the electrons shield each other from the nuclear charge. The net nuclear charge experienced by an electron is the effective nuclear charge, Zeff. If S is the average number of screening electrons: Zeff = Z - S

31. Atomic Size The radius of an atom is found from the distance between nuclei in a molecule.

32. Trends in Atomic Size Size increases going down a column of the periodic table. Size decreases from left to right in a row. There are two factors at work: principal quantum number, n, and the effective nuclear charge, Zeff.

33. Problem Solving

34. Ionization Energy The energy required to remove an electron from an atom in its ground level. Example: Hydrogen H(g) ? H+(g) + e? I = ?E

36. Trends in Ionization Energy The s electrons are more effective at shielding than p electrons. Therefore, forming the s2p0 becomes more favorable. When a second electron is placed in a p orbital, the electron-electron repulsion increases. When this electron is removed, the resulting s2p3 is more stable than the starting s2p4 configuration. Therefore, there is a decrease in ionization energy.

38. Electron Affinity The change in energy when an electron is added to a gaseous atom. Cl(g) + e? ? Cl?(g) ?E = -349 kJ/mol A large, negative ?E indicates strong attraction between the atom and the added electron. A positive ?E indicates the addition of an electron is unfavorable. Ne(g) + e? ? Ne?(g) ?E = 40 kJ/mol

42. Ionic vs. Atomic Radius