This lesson includes an Electron Configuration Worksheet, Guided Notes, Power Point Presentation, Exit Quiz, Vocabulary Activity, and much more! Start your lesson with this...

The Atomic Number Revisited:

The atomic number tells us the number of protons an atom possesses. In a neutral atom, the number of protons and electrons are equal. Therefore, the atomic number tells us the number of electrons an atom has. For example, Carbon has an atomic number of 6 and will, therefore, have six protons and six electrons.

Electron Arrangement:

The organization of the electrons in an atom is called its electron arrangement. Electrons occupy certain energy levels around the nucleus and only have a particular energy value, not a range of values. Electrons in atoms will occupy lower energy levels before they can occupy higher levels. The higher the energy level that an electron occupies, the more energy it possesses and the greater its distance from the nucleus. Each energy level has a limit to the number of electrons it can hold at any one time. The following shows the limits for the first twenty elements:

Electron Configuration in the First 20

• The first level holds a maximum of two electrons
• The second and third levels each hold a maximum of eight electrons
• The fourth level holds the remaining electrons.

Elements which have more than 20 protons are able to hold a maximum of 18 electrons in their third level.

Writing Electron Configurations:

The arrangement of electrons in the first twenty elements can be written using the rules above and commas to separate each energy level.  For example, carbon has 6 electrons, 2 of these electrons will occupy the first energy level and the other 4 will occupy the second level. Therefore, carbon is written as 2, 4.

Drawing Electron Configuration Diagrams

In some instances, you may be required to draw the structure of one of the first twenty atoms. In this case, knowing the electron configuration rules above for the first twenty elements and how to draw an atom is necessary. The diagrams below show the steps to drawing electron configuration diagrams.  Note these diagrams do not take into consideration the different shaped orbitals that atoms possess.

Valence Electrons

Electrons in the outermost (or highest) energy level are called valence electrons while those closer to the nucleus are called core electrons. Valence electrons determine the atoms stability and are involved in chemical reactions. Elements which have the same number of valence electrons in their outer shell tend to have similar chemical properties. For example lithium, sodium and potassium each have one lone valence electron, they all react violently when in contact with water to produce hydrogen gas.

The number of valence electrons an atom has dictates its behaviour in a chemical reaction and its reactivity. For example, sodium (Na) has one valence electron which is given away in a chemical reaction. By contrast, chlorine (Cl) has seven valence electrons and requires one more to fill its valence shell. Therefore, chlorine takes an electron in a chemical reaction.  Some atoms such as neon (Ne) and argon (Ar) have full valence shells are considered inert or stable. This means that they are do not give or take electrons in a chemical reaction.

Electron Orbitals: Electron Configuration Orbital Diagram Worksheet Answers

(The electron configuration orbital diagram worksheet answers can be found at the bottom of the lesson.)

The 2, 8, 8, 18 rule is a very simplistic view of electron configuration and doesn’t give the full picture when it comes to electron configuration. While it works for the first 20 elements, in order to look at other atoms higher than atomic number 20, we need to look closer at the types of orbitals in each electron shell in more detail.  Orbitals represent the space around the nucleus of an atom of which there is the greatest chance of locating an electron. Orbitals make up sub-shells, which make up electron shells (these are numbered 1-4).

The four different types of orbitals which are denoted by the letters s, p, d and f and correspond to the shape that electron wave takes in relation to its density being higher in some regions than others. The shapes of the four orbitals are pictured below.

The s-orbital is spherical in shape.  The p-orbital is a 3-dimensional dumb-bell shape. Its 3 orbitals exist in different planes (called x, y and z) and are right angles to one another. The d-orbital has four lobes, while the f-orbital is the most complex of the four has eight lobes.

The table below shows the types of orbitals present in each electron shell and the maximum number of electrons each type of orbital and shell can carry.

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Key Things to Note about Orbitals:

• There are only ever a maximum of 2 electrons per orbital.
• There is always an s-orbital in each electron shell.
• Whenever p-orbitals are present, there is a maximum of three p-orbitals that can be filled.
• Whenever d-orbitals are present there is a maximum of five d-orbitals that can be filled.

The arrangement of electrons must obey the following three rules:

1. The Aufbau Principle:

The orbitals with lower energies, closer to the nucleus are filled before those with higher energies.

This means that the 4s sub-shell will be filled before the 3d sub-shell as 4s has a slightly lower energy level, due its spherical shape allowing it to get closer to the nucleus. The 3d sub-shell cannot get as close to the nucleus so its energy level is higher.

Attribution: By CK-12 Foundation (raster), Adrignola (vector) - File:High School Chemistry.pdf, page 342, Public Domain, https://commons.wikimedia.org/w/index.php?curid=16749529

Electrons therefore fill orbitals in the following order:

1s, 2s, 2p, 3s, 3p, 4s, 3d, 4p, 5s, 4d, 5p, 6s, 4f, 5d, 6p, 7s, 5f.

2. The Pauli Exclusion Principle:

Electrons which occupy the same atomic orbital must spin in opposite directions; which therefore limits the number of electrons occupying an orbital to two.

3. Hund’s Rule of Maximum multiplicity:

Electrons occupy all the different orbitals within the same sub-level before doubling up inside orbitals.

The sequence below shows the order that p-orbitals are occupied for 1-6 electrons.

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Orbital Diagrams

Orbital diagrams are a pictorial way to describe the arrangement of the electrons in an atom. The orbitals are arranged from lowest to highest energy level, with arrows to indicate electrons with the opposite spin occupying the orbitals.

The example below shows you how to fill out these diagrams:

Nitrogen has 7 electrons its electron notation is 1s²2s²2p^3, so the orbitals are filled as follows:

Orbital Notation:

When writing out notation for elements there is a particular format that is used. It is shown below:

How to write electron orbital notation:

To write out the electron configuration you need to determine the number of electrons using the periodic table and fill up each orbital in order, starting with 1s, until you run out of electrons.

The steps are as follows:

• The first electron shell contains a single s-orbital, where the first two electrons are placed as this has the lowest energy level. This is denoted as 1s².
• The next two electrons are placed in the next lowest energy level found in the second electron shell called 2s². Further electrons (up to 6) can be placed in the 2 p-orbitals and are denoted as 2p⁶.
• From here, the third electron shell will have two electrons occupy its 3s orbital (3s²), followed by another 6 electrons in its 3p orbitals (3p⁶).
• Should a fourth electron shell be required another two electrons can be placed in the 4s orbital (4s²), before adding 10 electrons to the 3d orbital (3d¹⁰) as this has a lower energy level than the 4p orbital.

The Diagonal Rule

The diagonal rule can be used to determine an electronic configuration.

Shortened Notation

The following steps can be used to write abbreviated electron configurations.

1. Find the symbol for the element on a periodic table.
2. Write the symbol in square brackets for the noble gas located at the far right on the horizontal row before the element.
3. Return to the row containing the element you wish to describe and to the far left. Following the elements in the row from left to right, write the outer-electron configuration associated with each column until you reach the element you are describing.

Notation for Ions

Ions are atoms which have lost or gained electrons. The mechanism by which this occurs is explained in later lessons. Therefore electron configurations for ions need to have electrons added or subtracted from the total.  If an ion has a positive (+) charge electrons must be subtracted from the total. Ions, with a negative charge, require electrons to be added to the total.

Exceptions to the rule

Chromium and copper are the two main exceptions to rules for electron configurations. In these cases, a completely full (3d¹⁰) or half full (3d⁵) d sub-level is more stable than a partially-filled d sub-level, so an electron from the 4s orbital is excited and rises to one of the 3d orbitals to fill it.

• The transition metal chromium and has 24 electrons. The electron notation for chromium is 1s², 2s², 2p⁶, 3s², 3p⁶, 4s², 3d⁴.

The d sublevel is 1 electron short of it being half filled. So one of the electrons from the 4s orbital is moved over to the 3d orbital.

Instead chromium becomes 1s², 2s², 2p⁶, 3s², 3p⁶, 4s¹, 3d⁵ or [Ar] 4s¹, 3d⁵ using shortened notation. This makes a half-filled d-orbital which is much more stable than d⁴.

• The other exception is copper with a notation of 1s², 2s², 2p⁶, 3s², 3p⁶, 4s², 3d⁹. Once again, an electron from the 4s is moved over to the 3d orbital completely filling the d sublevel making it stable. So, it becomes 1s², 2s², 2p⁶, 3s², 3p⁶, 4s¹, 3d¹⁰ or [Ar] 4s¹, 3d¹⁰.

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Unit 1 – Structure and Properties of Matter

• 1-1 Atoms and Molecules

• 1-2 The Kinetic Molecular Theory and Properties of Matter

• 1-3 Models of an Atom

• 1-4 Atomic Structure

• 1-5 Isotopes

• 1-6 Electron Configuration

• 1-7 The Periodic Table of Elements