Electron Configuration and Atomic Properties

1. Electron Configuration and Atomic Properties Topics: Electron Spins and Magnetism Orbital Energy Electronic Configurations of Elements Atomic Properties Ions OWLs Due on 28-November.

2. Electron Spin and Magnetism • There is an additional quantum number, and one that is very important! • Spin Quantum Number (ms) • Electrons can have a spin of +1/2 or -1/2 • Since electrons are charged particles, as they move (spin, for example), they create magnetic fields.

3. We’ll see how we fill electrons with spins into orbitals soon, but we need to know that there are 3 types of magnetic materials: • Diamagnetic (non-magnetic) • Paramagnetic (weakly magnetic) • Ferromagnetic (strongly magnetic)

4. Orbital Energies (single e- species) • In single electron species (hydrogen) or even some ions, the orbitals at each energy level have the same energy. • Even if the orbitals are different in physical size • This is a very simplistic model that works for few species

5. Orbital Energies (multiple e- species) • In multiple electron species, orbital energies are intermixed among different “n” energy levels. • Most common • Leads to electron configurations and systematic filling for most elements and ions. • The energies change because electrons now interact with other electrons • Repulsive forces

6. Electron Configuration in Atoms • The Pauli Exclusion Principle states, simply, that no two electrons may have the same set of quantum numbers. • n • l • ml • ms • Additionally, atomic orbitals are filled from the lowest energy up when the atom is in the “ground” state • Lowest energy state

7. Electron Configuration in Atoms • This results in the following order: 1s, 2s, 2p, 3s, 3p, 4s, 3d, 4p, 5s, 4d, 5p, 6s, 4f, 5d, 6p, 7s, 5f, 6d, 7p • You should know up through 4f • The total number of electrons is the same as the atomic number • We have as many unpaired electrons in a specific orbital (s, p, d, f) as we can fit. • Hund’s Rule • And what does this mean for the Periodic Table? • Flashback to the previous chapter……

9. Let’s Build Some Electron Configurations! • Remember that there are always a few exceptions • Focus on the rules rather than the exceptions. • Remember your filling order. • Fill in a box diagram • Fill in an energy level diagram • Write it in notation • Use the atomic number and position of the element on the periodic table to help!!!!

10. Hydrogen (ground state): • Atomic Number 1 • # of electrons = 1 • n = 1 • l = 0 (s orbital) • ml = 0 (one s orbital) • ms = +1/2 • 1s1

11. Lithium (ground state) • Atomic Number 3 • # of electrons = 3 • n = 1, 2 • l = 0,1 (s and p orbitals) • ml = 0 (one s orbital) • ml= -1, 0, 1 (3 p orbitals) • ms = +1/2 and -1/2 • 1s2 2s1 • Helium (ground state) • Atomic Number 2 • # of electrons = 2 • n = 1 • l = 0 (s orbital) • ml = 0 (one s orbital) • ms = +1/2 and -1/2 • 1s2

12. Carbon (ground state) • Atomic Number 6 • # of electrons = 6 • n = 1, 2 (second period on table) • l = 0, 1 (s and p orbitals) • ml = 0 (one s orbital), -1,0,1 (3 p orbitals) • ms = +1/2 and -1/2 • 1s22s22p2

13. What element is this? What is the electronic configuration for Calcium?

14. Shorthand Notation…… • Sometimes we abbreviate electron configurations using Noble Gas Notation: • What element is this? • [Ar]3d104s24p5

16. Scandium (Sc)

17. Back to the Periodic Table….. Mendeleev’s periodic table

20. There are other types of electron configurations we need to consider (and terms) • Ground (lowest) versus excited (higher) energy state…. • Inner (core) shell and valence (outer) shell electrons • Atoms versus ions (a topic for later in Chapter 7)

21. Ground vs. Excited State • An excited state electronic configuration is present when an atom (or ion) absorbs energy and an electron is promoted to a higher energy level • This can occur even if the n energy level does not exist in the ground state. The level is still there. • Ground State Mg (1s22s22p63s2), 12 electrons. • Excited State Mg (just one of many possible) • 1s22s22p63s13p1 • The atom must absorb energy for this to happen. • When it transitions back to the ground state, that exact amount of energy is given off • Light • Heat • Kinetic energy transferred to another atom or molecule.

22. Inner versus Outer Shell • Inner (core) shell electrons are those in full (“closed”) n energy levels • Ordinarily those seen in the noble gas configuration. • Take Aluminum for example • 1s22s22p63s23p1 • [Ne] 3s23p1 • [Ne] electrons represent the inner or core shell electrons • The 3s2 and 3p1 electrons are the valence or outer shell. • When ions are formed, only valence electrons are gained or lost • Al3+

23. Periodic Trends and Properties • Effective Nuclear Charge (Z*) • As atomic number increases, so does the number of protons. The nuclear charge increases, which raises the energy of orbitals surrounding the nucleus. • Effective nuclear charge (Z*) = Z – (# of core shell electrons) • It is a relative number, used for basic comparisons. • It represents the nuclear charge experienced by the highest energy, valence electrons

25. Atomic Size • Closely related to orbital configuration, energies and effective nuclear charge • Atoms with a greater nuclear charge “pull” electrons in closer to the nucleus and are smaller • Opposites attract

26. Covalent radius: Distance between two nuclei when two atoms are bonded together (Cl2 as an example) Metallic radius: Distance between nuclei when atoms of a metallic element are near each other in a metallic crystal (say a block of Zn)

27. Ionization Energy • The energy required to remove an electron from an element in a gaseous state. • Increases across a period because of increasing orbital energy and effective nuclear charge

28. Ionization energy trends

29. Ionization energy trends

30. Electron Affinity • The energy change when a gaseous atom gains an electron • More negative values represent a greater affinity.

32. Ions… • Filled shells (n energy levels) or orbitals, represent the most stable electron configurations • Electrons of opposite spin are paired • Energy levels may be full • Ions form because the energy level of that electron configuration is particularly stable • Most ions represent elements trying to achieve noble gas electron configurations • Valence (outer) shell electrons are lost (cations) or gained (anions) to produce ions.

33. Cations (atom loses electrons)

34. Anions (atom gains electrons)

35. Ion sizes • Anions are normally larger than their atom “parent”, because they are gaining electrons • Cations are normally smaller than their atom “parent” because they are losing electrons • Anions are generally larger than cations