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2016 Topic 2: Electron Configuration

This document discusses electron configuration and how it is described using quantum numbers. It explains the main concepts like energy shells (n), subshells (s, p, d, f), orbitals and how electrons fill these according to the Aufbau principle and Hund's rule. Examples are provided to show the electron configurations of elements like hydrogen, helium, lithium and how the configurations change as the atomic number increases. Practice problems are included at the end to determine configurations of additional elements.

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Electron Configuration IB Chemistry Power Points Topic 2 Atomic Structure www.pedagogics.ca
Going beyond 2,8,8,2 Electron configuration in atoms can be described by terms called quantum numbers – no two electrons can have the same quantum number! st 1 Term: shell (n) - main energy level n = 2 n = 1 n = 3 lone electron of hydrogen energy Issue with this graphic?
2 nd Term: subshell - designated by s, p, d, f - designates the sub-energy level within the shell. - refers to the shape(s) of the volume of space where n = 3 n = 2 1s electrons are be located. The first shell (1) has one subshell (s). The s subshell has 1 spherical shaped orbital orbitals are volumes of space where the probability of finding an electron is high energy
The Electronic Configuration of Hydrogen 1s Hydrogen has one electron located in the first shell. (Aufbau principle – electrons will occupy the lowest energy orbitals first) The first shell has only one subshell (s). The s subshell holds a single s orbital. Electronic configuration 1 1s shell subshell # of electrons present energy 1s  Orbital Energy Level Diagram
The Electronic Configuration of Helium He: Atomic # of 2, 2 electrons in a neutral He atom 1 H 1s 2 He 1s 1s He 1s  energy the maximum number of electrons in an orbital is TWO if there are 2 electrons in the same orbital they must have an opposite spin. This is called Pauli’s Exclusion Principle
Lithium (Li) Li: Z=3 Li has 3 electrons. 1s 2 nd shell 1s The 2nd shell (n= 2) has 2 subshells which are s and p. The s subshell fills first! (Aufbau Principle) 2s 2p 2s  Li 1s  Orbital Energy Level Diagram 2 2s Li 1s 1 Electronic configuration energy
Subshells so far - designated by s, and p - refers to the shape(s) of the volume in which the electron can be located. - also designates an energy level within the shell. - relative energy: s < p s subshell: spherical 1 orbital x y z x y z p subshell: pair of lobes, 3 orbitals, each holds 2 electrons
Berylium (Be) Be: Z=4 Be has 4 electrons. Be 1s 2 2s 2 2s   Be 1s  Electronic configuration Orbital Energy Level Diagram Boron (B) has 5 electrons, the s subshell is full so the 5 1s 2 nd shell 2s 2p B 1s 2 2s 2 2p 1 2p  2s  B 1s  th electron occupies the first orbital in the p subshell energy
Carbon (C) C: Z=6 C has 6 electrons. 1s 2 nd shell 2s 2p C 1s 2 2s 2 2px 1 py 1 2p   2s  C 1s  C 1s 2 2s 2 2p 2 The 6 th electron occupies an empty p orbital. This illustrates “Hund’s Rule” – electrons do not pair in orbitals until each orbital is occupied with a single electron. The electron configuration is But always written as
2p    2s  N 1s  1s 2 2s 2 2p 3 2p    2s  O 1s  1s 2 2s 2 2p 4 2p    2s  Ne 1s  1s 2 2s 2 2p 6
Practice Use the sheets provided to fill out orbital diagrams and determine the electron configuration for the following elements 1. Fluorine 2. 56Fe 3. Magnesium - 22 4. 131I 5. Potassium – 42 6. 75Ge 7. Zirconium – 90 8. 41Ca2+
Practice Use the sheets provided to fill out orbital diagrams and determine the electron configuration for the following elements 1. Fluorine 1s 2 2s 2 p 5 2. 56 Fe 1s 2 2s 2 p 6 3s 2 3p 6 4s 2 3d 6 3. Magnesium – 22 1s 2 2s 2 p 6 3s 2 4. 131 I 1s 2 2s 2 p 6 3s 2 3p 6 3d 10 4s 2 4p 6 4d 10 5s 2 5p 5 5. Potassium – 42 1s 2 2s 2 p 6 3s 2 3p 6 4s 1 6. 75 Ge 1s 2 2s 2 p 6 3s 2 3p 6 4s 2 3d 10 4p 2 7. Zirconium – 90 1s 2 2s 2 p 6 3s 2 3p 6 4s 2 3d 10 4p 6 5s 2 4d 2 8. 41 Ca 2+ 1s 2 2s 2 p 6 3s 2 3p 6
Electron Configurations and the Periodic Table So far, we have seen how the subshell model provides and explanation for the patterns in ionization energy we see in the periodic table. You have also seen how to write electron configurations Example CALCIUM  1s 2 2s 2 2p 6 3s 2 3p 6 4s 2 - Principle energy level subshell # of e Calcium can also be written shorthand as: 2 [Ar]4s
The organization of the Periodic table correlates directly to electron structure
Condensed electron configurations – for example the electron 2 10 5 configuration of bromine can be written [Ar] 4s 3d 4p Read questions carefully – many IB questions require you to write the FULL electron configuration
You are responsible for configurations up to Z = 36 (Kr). The table works well for this with the exception of Cr and Cu
Chromium’s configuration is: 1 3d [Ar]4s 5 Copper’s configuration is: [Ar]4s 1 3d 10 These configurations are energetically more stable than the expected arrangements. KNOW THEM!
Electron configuration of ions: In general, electrons will be removed from orbitals (ionization) in the reverse order that the orbitals were filled. In other words, electrons vacate higher energy orbitals first. The exception: TRANSITION METAL IONS When these ions form, electrons are removed from the valence shell s orbitals before they are removed from valence d orbitals when transition metals are ionized. For example: Cobalt has the configuration [Ar] 4s 2 3d 7 The Co 2+ 3+ and Co ions have the following electron configurations. Co 2+ : [Ar] 3d 7 Co 3+ : [Ar] 3d 6
Condensed electron configurations – for example the electron 2 10 5 configuration of bromine can be written [Ar] 4s 3d 4p 1. Si ___________________________ 2. S2- ___________________________ 3. Rb+ ___________________________ 4. Se ___________________________ 5. Ar ___________________________ 6. Nb ___________________________ 7. Zn2+ ___________________________ 8. Cd ___________________________ 9. Sb ___________________________
Review: the principles involved Aufbau Principle: electrons will fill the lowest energy orbitals first Hund’s Rule: the most stable arrangement of electrons in orbitals of equal energy is where there is the maximum number of unpaired electrons all with the same spin. Pauli’s Exclusion Principle: A maximum of two electrons can occupy a single orbital. These electrons will have opposite spins.