Electron Configurations Electrons can be distributed amongst the subshells/orbitals of an atom in different ways –producing different electron configurations. The most stable configuration has the lowest energy – corresponding to the situation where electrons get as close to the nucleus as possible while staying as far away from each other as possible.

Electron Configurations Chemistry 140 Fall 2002 Electron Configurations Aufbau process Electrons occupy orbitals in a way that minimizes the energy of the atom. Pauli exclusion principle No two electrons can have all four quantum numbers alike. Hund’s rule When orbitals of identical energy (degenerate orbitals) are available, electrons initially occupy these orbitals singly. Figure 8-36 suggests the order in which electrons occupy the subshells in the principal electronic shells; first the 1s then 2s, 2p and so on. The exact order of filling of orbitals has been established by experiment. In 1926, Wolfgang Pauli explained complex features of emission spectra associated with atoms in magnetic fields by proposing that no two electrons in an atom can have all four quantum numbers alike. The first three quantum numbers, n, l, and ml determine a specific orbital. Two electrons may have these three quantum numbers alike; but if they do, they must have different values of the spin quantum number. Another way to state this result is that only two electrons may occupy the same orbital, and these electrons must have opposing spins. Copyright © 2011 Pearson Canada Inc. General Chemistry: Chapter 8

Populating Orbitals If we continued with the previous exercise we’d see that for each value of n there is one s orbital (unique ml value), for n ≥ 2 three p orbitals (three ml values) and, for n ≥ 3, five d orbitals (five ml values: -2, -1, 0, +1, +2). This means that s, p, d subshells can contain (at most) 2, 6 and 10 electrons respectively. Relative subshell energies comes from experiment – next slide.

The order of filling of electronic subshells FIGURE -37 The order of filling of electronic subshells Copyright © 2011 Pearson Canada Inc. General Chemistry: Chapter 8

Electron Configurations - Monatomic Species We will usually write condensed electron configurations which show the number of electrons per subshell (usually for the ground state or lowest energy configuration possible). Excited states or higher energy configurations will also be considered. We will also write expanded electron configurations where –particularly for valence electrons – the number of electrons per orbital is shown.

“Condensed” Electron Configurations Example of C Atom – 6 Electrons Subshell Maximum e-s Per Subshell Actual # e-s Number of e-s Accommodated “Left Over” 1s 2 4 2s 2p 6 3s 3p

Electron Configurations – “Exponential” Notation The information contained in the previous slide can be represented more compactly using a notation where “exponents” are used to indicate the number of electrons per subshell. For the neutral C atom in its ground electronic state the electron configuration written in exponential notation is 1s22s22p2.

“Expanded” Electron Configurations A more detailed picture of electron the electronic “structure” of an atom or monatomic ion is obtained using so-called “expanded” configurations. The expanded configurations reflect Hund’s Rule (based on experiment) – electrons occupy equivalent orbitals to the maximum extent possible and with their spins parallel.

“Expanded” Electron Configurations Example of C Atom – 6 Electrons Subshell Maximum e-s Per Subshell Actual # e-s Number of e-s Accommodated “Left Over” 1s 2 4 2s 2px 1 5 2py 6 2pz

“Expanded” Electron Configurations For a C atom in its ground state the expanded electron configuration could be written as 1s22s22px12py1 OR 1s22s22px12pz1 OR 1s22s22py12pz1. All configurations have the same energy. Expanded configurations help us identify any unpaired electrons that might be present. Unpaired electrons are important since they can give rise to important magnetic properties (paramagnetism).

Expanded Electron Configurations & Orbital Diagrams The detailed electron configurations of atoms and monatomic ions can also be represented using orbital diagrams. Here a box represents each orbital and arrows indicate both the “spins” of electrons (ms value +ve or –ve) and whether the orbital contains 0, 1 or 2 electrons. The next slide represents the orbital diagram for the C atom. Note the “parallel spins” for the two 2p electrons.

Representing Electron Configurations spdf notation (condensed) 1s22s22p2 spdf notation (expanded) 1s22s22px12py1 spdf notation Copyright © 2011 Pearson Canada Inc. General Chemistry: Chapter 8

General Chemistry: Chapter 8 Chemistry 140 Fall 2002 The Aufbau process To write electron configurations we will use the aufbau process. Aufbau is a German word that means “building up,” and what we do is assign electron configurations to the elements in order of increasing atomic number. To proceed from one atom to the next, we add a proton and some neutrons to the nucleus and then describe the orbital into which the added electron goes. Copyright © 2011 Pearson Canada Inc. General Chemistry: Chapter 8

Transition Metals – Electron Configurations Transition metals have surprisingly rich chemistry. The electronic configurations of the neutral atoms are relatively complex since d subshells (with 5 orbitals) are being filled. For the first series of transition metals the 4s and 3d subshells have similar energies and surprises are seen for their electronic configurations. Cr and Cu do not have the “expected” electron configurations. Copyright © 2011 Pearson Canada Inc. General Chemistry: Chapter 8

The Aufbau Process – Sc through Zn Copyright © 2011 Pearson Canada Inc. General Chemistry: Chapter 8

Transition Metal Configurations The unexpected electron configurations found experimentally for Cu and Cr are often rationalized in terms of a special stability (low energy) associated with a half full (3d5) and full (3d10) d subshell. Similar issues arise with transition metal ions. The large numbers of unpaired electrons seen for transition metals gives them interesting magnetic properties. Transition metal compounds are often colourful – discussed in higher level courses.

Electron Configurations and the Periodic Table Chemistry 140 Fall 2002 Electron Configurations and the Periodic Table s block. The s orbital of highest principal quantum number (n) fills. The s block consists of groups 1 and 2 (plus He in group 18). p block. The p orbitals of highest quantum number (n) fill. The p block consists of groups 13, 14, 15, 16, 17, and 18 (except He). d block. The d orbitals of the electronic shell n-1 (the next to outermost) fill. The d block includes groups 3, 4, 5, 6, 7, 8, 9, 10, 11, and 12. f block. The f orbitals of the electronic shell n-2 fill. The f-block elements are the lanthanides and the actinides. FIGURE 8-38 Copyright © 2011 Pearson Canada Inc. General Chemistry: Chapter 8

Valence Shell Configurations The occupied shell with the highest value of n is called the valence shell. When atoms undergo chemical change electrons in the valence shell can be lost or shared with other atoms. The valence shell can also pick up electrons. Atoms with similar chemical properties often have the “same” valence shell electron configuration. For example, Li, Na, K, Rb, Cs and Fr have an ns1 valence shell configuration.

General Chemistry: Chapter 8 Chemistry 140 Fall 2002 Copyright © 2011 Pearson Canada Inc. General Chemistry: Chapter 8

Electron Configurations - Examples 1. Write condensed electron configurations representing the ground electronic states of the P atom and the P3- ion. 2. Write condensed electron configurations representing the ground electronic states of the Sr atom and the Sr2+ ion. 3. Write condensed electron configurations for the ground and two excited electronic states of the Na atom.

Electron Configurations - Examples 4. Construct orbital diagrams for (a) the Al atom, (b) the Si atom (for Bill Gates), (c) the S atom and (d) the S2- ion. 5. How many unpaired electrons are there in a neutral arsenic (As) atom? 6. Write a set of four possible quantum number values (n, l, ml and ms) for an electron in the (a) valence shell of Mg and (b) the highest occupied orbital (energy) of Pb.

Electron Configurations - Examples 7. Write chemical symbols for five monatomic species having the ground state electron configuration 1s22s22p63s23p6. How many electrons are protons are contained in each species? 8. How many unpaired electrons do neutral Al, Si, P, S, Cl and Ar atoms possess?

Final Note on Nodes The H atom wave functions tell us that there are only radial nodes for the 1s, 2s, 3s.. orbitals. Angular or planar nodes become important as we move to p orbitals (one planar node) and d orbitals (2 planar nodes). This is seen in the slide on the next table and graphical representations of d orbitals. (Orbital shapes impt in chemical bonding.)

Hydrogen Atom Wavefunctions – Number of Nodes Orbital Designation Total # Nodes (n-1) Planar Nodes (l) Radial Nodes 1s 0 (s orbital) 2s 1 2p 1 (p orbital) 3s 2 3p 3d 2 (d orbital)

Representations of the five d orbitals FIGURE 8-30 Representations of the five d orbitals Copyright © 2011 Pearson Canada Inc. General Chemistry: Chapter 8

The Periodic Table In studying the electronic structure of atoms we mentioned that chemical families of elements have similar valence shell electron configurations. Historically, however, a detailed knowledge of atomic structure came subsequent to the observation that groups of elements had similar chemical properties.

The Periodic Table In the modern Periodic Table elements are arranged in order of increasing atomic number so that groups of elements with similar chemical properties appear in columns. The 100+ known elements are commonly divided into three groups – the metals, non-metals and the metalloids. These three sets of elements have rather different physical properties.

Metals and Nonmetals and Their Ions Good conductors of heat and electricity. Malleable and ductile. Moderate to high melting points. Nonmetals Nonconductors of heat and electricity. Brittle solids. Some are gases at room temperature. Metalloids Metallic and non-metallic properties Slide 28 of 35 Copyright © 2011 Pearson Canada Inc. General Chemistry: Chapter 9