Lesson 3: Electrons

Part a: Electron Configurations

Part 3a: Electron Configurations
Part 3b: Electrons and the Periodic Table
Part 3c: Exceptions to the Rule

 

Electrons and Chemical Properties

We discussed chemical properties in Chapter 2. Chemical properties describe how a substance interacts with other substances to form compounds. The chemical properties of an element include descriptions of how that element combines with other elements and the ratio of atoms in the resulting compound. These chemical properties are determined by the manner in which electrons are arranged in the atoms of those elements.
 
The quantum mechanical model will be used in Lesson 3 to build a foundation for understanding chemical properties. The foundation relies on the use of electron configurations. Electron configurations are a symbolic means of showing the location of electrons within an atom. Electrons are housed inside of orbitals. An electron configuration identifies the address of all the houses and the number of electrons residing in the houses. The orbitals are identified by standard orbital notation (1s, 2s, 2p, 3d, etc.) and a superscript indicates how many electrons are housed inside orbitals of that type.
 
 

 

Energy Levels

To write an electron configuration for an element, one must know the names of the orbitals and their relative energies. This was discussed quite thoroughly in Lesson 2. If the following brief review is not sufficient, please review Lesson 2c. The graphic at the right identifies the names (orbital notation) of all the orbitals and places them on a graph to show their relative energy. Generally, orbitals with a smaller n value (the number in the orbital notation) have lower energies. As was discussed in Lesson 2c, there are some surprises for any principal energy level containing d orbitals. The s orbitals of the 4^th energy level are lower in energy than the d orbitals of the 3^rd energy level. The same can be said for all other energy levels with an even higher n. The graph also shows that the 5s, 5p, and 6s orbitals are lower in energy than the 4f orbitals.
 
These surprise orderings make the task of ordering the orbitals by energy very memory intensive. One means of generating the order involves the use of the diagonal rule (sometimes referred to as the Madelung Rule). The orbitals are listed by principal energy level. Then diagonals are drawn through the listing. The use of this rule is depicted in the diagram. More details can be found in Lesson 2c.
 

 

 
 

 

 

 

Electrons Filling Orbitals

Writing electron configurations for elements also demands an understanding of the order by which electrons enter the orbitals. This was also thoroughly discussed in Lesson 2c so the discussion here will be brief. There are three rules to adhere to when deciding the orbital that the 6^th, the 8^th, the 14^th, the 16^th, etc. electron will be housed in.

1. Aufbau Principle: electrons will completely fill the lowest energy orbitals first before entering orbitals of the next highest energy.
2. Hund’s Rule: electrons will half-fill the orbitals of a given energy sublevel with the same spin direction before pairing up inside such orbitals.
3. Pauli Exclusion Principle: when electrons pair up inside of orbitals, they do so with opposite spin direction.

 
These rules and their application were thoroughly demonstrated in Lesson 2c. Orbital box diagrams were used to show the placement of electrons. Electrons are represented by arrows. The direction that they point (up or down) reflect their spin direction. Examples of orbital box diagrams discussed in Lesson 2c are shown here:
   
 
 

Orbital Box Diagrams to Electron Configurations

Orbital box diagrams are useful tools for determining electron configurations. Once the diagram is completed, begin writing the electron configuration. Begin on the left side of the diagram with the lowest energy orbital. Write the orbital notation of the orbital (1s). Add a superscript to indicate the number of electrons in the orbital (1s²). Continue from left (lowest energy) to right (higher energies) across the orbital box diagram. Write the orbital notation plus superscript for any orbital type with electrons. The process is shown for carbon and chlorine below.
   

   

Examples of Electron Configurations

Eleven more examples of electron configurations are shown below. Observe how the orbital box diagram naturally translates into an electron configuration.
   
 
 

Abbreviated Electron Configurations

Electron configurations can become quite long. Consider the configuration for barium with 56 electrons:
   
After writing a few barium-like configurations, one begins thinking is there a shortcut? The good news is that there is. The shortcut is to write the so-called abbreviated electron configuration. Writing the abbreviated form involves the following:
 

1. Locate the element on the periodic table.
2. Identify the noble gas element (Group 18) at the end of the row above the element.
3. Write the symbol of the noble gas enclosed in brackets. That takes care of all the electrons included in the noble gas. For instance, if you write [Xe], you are accounting for the first 54 electrons.
4. Write the remainder of the electron configuration, beginning with the electron after the noble gas’s last electron. For the abbreviated electron configuration of barium, write [Xe]. Then add the location of the 55^th and 56^th electrons.

 
To use the shortened configuration, one must know either the orbital notation for the element at the end of the row or the orbital notation for the element at the beginning of a row. This is shown below. (We will have more to say about it in Lesson 3b.)
   
The table below shows the abbreviated electron configuration for a variety of elements. Notice that they always begin with the symbol of a noble gas element enclosed in brackets. Once that is written, the energy level diagram showing the ordering of orbitals is used to finish out the configuration in the usual manner.
     

Electron Configuration for Ions

The topic of ions was discussed in Chapter 3. Ions have unequal numbers of protons and electrons and are thus charged. An ion can form when an atom gains or loses one or more electrons. Electron configurations can be written for charged ions in the same way they are written for neutral atoms. The first step is to determine the number of electrons in the ion. Look up the atomic number on the periodic table. This indicates the number of protons. Use the atomic number and the charge of the ion to determine the number of electrons. A positively charged ion has less electrons than protons. A negatively charged ion has more electrons than protons. The table shows how these principles can be used to determine the number of electrons in an ion.
   
Once the number of electrons is determined, enter the electrons into the orbital box diagram. Then translate the diagram into an electron configuration. Examples are shown below. Note how the electron configuration of a charged ion with 18 electrons is the same as the electron configuration of a neutral atom with 18 electrons.
   
 
 

Electron Shells

The electrons in atoms are located in orbitals. The collection of orbitals at a given energy level (n value) form an electron shell. For instance, the one 2s orbital and the three 2p orbitals of form the n=2 electron shell.
   
In Lesson 3 and subsequent chapters of this Chemistry Tutorial, we will distinguish between the outer shell electrons and the core electrons. Outer shell electrons, sometimes referred to as valence shell electrons, are s- and p-orbital electrons in the outermost electron shell of atoms. Valence shell electrons are the electrons that are involved in bonding. And because of this, they are the electrons that determine the chemical properties of elements. Whether sodium and oxygen form NaO or NaO₂ or Na₂O or NaO₃ or Na₃O is dependent upon the number of valence shell electrons in the atoms of the two elements. The core electrons are those electrons that are in the inner shells of atoms; they typically do not involve themselves in bonding. The number of valence shell electrons is evident when you inspect an electron configuration.
   
Note that d-orbital electrons are core electrons and are not counted towards the total number of valence electrons. Only s- and p-orbital electrons are counted towards the total.
 
 
 

Electron Shell Diagrams

The core and valence shell electrons are often represented by electron shell diagrams. An electron shell diagram shows the electrons layered in shells around a nucleus. Each principal energy level becomes an electron shell and the electrons of that level are represented by an X or a circle on the diagram. The electron shell diagram for oxygen and magnesium are shown below.
   
The notation “2, 6” and  “2, 8, 2” shown below the diagram is commonly added. The numbers indicate the number of electrons in each shell, beginning with the innermost electron shell. The last number represents the outer shell or valence shell electrons. There are 2 core electrons and 6 valence shell electrons in oxygen. There are a total of 10 core electrons and 2 valence shell electrons in magnesium.
 
 
 
In the next part of Lesson 3 we will discover one of the slickest tricks in Chemistry. The electron configurations of elements are related to the location of the element in the periodic table. This relationship allows us to make quick work of writing electron configurations in complete and abbreviated forms.
 
 
 
 

Before You Leave

• Practice. Try our Complete Electron Configurations Concept Builder. It’s great practice!
• Have fun while learning. Try our Periodic Table Battleship Concept Builder. It’s great practice and lots of fun!
• Download our Study Card on Electron Configurations. Save it to a safe location and use it as a review tool.
• The Check Your Understanding section below include questions with answers and explanations. It provides a great chance to self-assess your understanding.

 
 
 

Check Your Understanding

Use the following questions to assess your understanding. Tap the Check Answer buttons when ready.
 
1. Construct orbital box diagrams for the following elements:

1. ₁₅P  
   
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   Answer: 
   
   
   
    
   
    
2. ₂₆Fe  
   
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   Answer: 
   
   
    
   
    
3. ₃₅Br  
   
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   Answer: 
   
   
   
    
   
    
4. ₄₀Zr  
   
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   Answer: 
   
   
   
    
   
    

 
 
2. Use orbital notation like 3s or 3p to identify the orbital into which the following electrons will be placed. And indicate if it will immediately join an earlier placed electron to be a paired electron or if it enters as an unpaired electron.

1. 7^th electron 
   
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   Answer: 2p orbital, unpaired
   
    
   
    
2. 14^th electron 
   
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   Answer: 3p orbital, unpaired
   
    
   
    
3. 20^th electron 
   
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   Answer: 4s orbital, paired
   
    
   
    
4. 30^th electron 
   
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   Answer: 3d orbital, paired
   
    
   
    
5. 32^nd electron 
   
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   Answer: 4p orbital, unpaired
   
    
   
    
6. 52^nd electron 
   
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   Answer: 5p orbital, paired
   
    
   
    

 
 
3. Write the complete (not abbreviated) electron configuration for the following elements:

1. ₄Be 
   
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   Answer: 1s² 2s² 
   
    
   
    
2. ₉F 
   
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   Answer: 1s² 2s² 2p⁵ 
   
    
   
    
3. ₁₂Mg 
   
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   Answer: 1s² 2s² 2p⁶ 3s² 
   
    
   
    
4. ₁₄Si 
   
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   Answer: 1s² 2s² 2p⁶ 3s² 3p²   
   
    
   
    
5. ₂₀Ca 
   
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   Answer: 1s² 2s² 2p⁶ 3s² 3p⁶ 4s²  
   
    
   
    
6. ₂₇Co 
   
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   Answer: 1s² 2s² 2p⁶ 3s² 3p⁶ 4s² 3d⁷ 
   
    
   
    
7. ₃₄Se 
   
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   Answer: 1s² 2s² 2p⁶ 3s² 3p⁶ 4s² 3d¹⁰ 4p⁴  
   
    
   
    
8. ₄₁Nb 
   
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   Answer: 1s² 2s² 2p⁶ 3s² 3p⁶ 4s² 3d¹⁰ 4p⁶ 5s² 4d³  
   
    
   
    
9. ₅₀Sn 
   
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   Answer: 1s² 2s² 2p⁶ 3s² 3p⁶ 4s² 3d¹⁰ 4p⁶ 5s² 4d¹⁰ 5p² 
   
    
   
    

 
 
4. Write the abbreviated electron configurations for the following elements:

1. ₈O 
   
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   Answer: [He] 2s² 2p⁴ 
   
    
   
    
2. ₁₁Na 
   
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   Answer: [Ne] 3s¹ 
   
    
   
    
3. ₁₇Cl 
   
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   Answer: [Ne] 3s² 3p⁵ 
   
    
   
    
4.  ₁₈Ar 
   
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   Answer: [Ne] 3s² 3p⁶ 
   
    
   
    
5. ₃₅Br 
   
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   Answer: [Ar] 4s² 3d¹⁰ 4p⁵ 
   
    
   
    
6. ₄₄Ru 
   
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   Answer: [Kr] 5s² 4d⁶ 
   
    
   
    
7. ₅₂Te 
   
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   Answer: [Kr] 5s² 4d¹⁰ 5p⁴ 
   
    
   
    
8. ₆₃Eu 
   
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   Answer: [Xe] 6s² 4f⁷ 
   
    
   
    
9. ₇₂Hf 
   
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   Answer: [Xe] 6s² 4f¹⁴ 5d²  
   
    
   
    
10. ₈₅At 
    
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    Answer: [Xe] 6s² 4f¹⁴ 5d¹⁰ 6p⁵ 
    
     
    
     

 
 
5. Write the electron configuration for the following ions:

1. O^2- 
    
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   Answer: 1s² 2s² 2p⁶  
   
    
   
    
2. Ga³⁺ 
    
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   Answer: 1s² 2s² 2p⁶ 3s² 3p⁶ 4s² 3d⁸ 
   
   (The above configuration would be expected based on our rules and our energy levels for orbitals. For reasons not discussed until later, the configuration may more likely be: 1s² 2s² 2p⁶ 3s² 3p⁶ 3d¹⁰ with no 4s electrons.)
   
    
   
    
3. Mg²⁺  
   
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   Answer: 1s² 2s² 2p⁶  
   
    
   
    
4. Ti²⁺  
   
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   Answer: 1s² 2s² 2p⁶ 3s² 3p⁶ 4s²  
   
    
   
    
5. Se^2-  
   
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   Answer: 1s² 2s² 2p⁶ 3s² 3p⁶ 4s² 3d¹⁰ 4p⁶   
   
    
   
    
6. I^-  
   
   See Answer
   Answer: 1s² 2s² 2p⁶ 3s² 3p⁶ 4s² 3d¹⁰ 4p⁶ 5s² 4d¹⁰ 5p⁶  
   
    
   
    

 
 
6. Construct an electron shell diagram for the sodium atom. Identify the number of core electrons and valence shell electrons.



 
7. Construct an electron shell diagram for the sulfur atom. Identify the number of core electrons and valence shell electrons.




8. If you were to draw the electron shell diagram for the following ions: O^2-, F^-, Na⁺, and Mg²⁺, then what would they all have in common?




9. Determine the number of valence shell electrons in the following atoms:

1. ₁₂Mg 
   
   See Answer
   Answer: 2
   
   There are two electrons in its 3s orbital.
   
    
   
    
2. ₁₅P 
   
   See Answer
   Answer: 5
   
   There are two electrons in its 3s orbital and three electrons in its 3p orbitals.
   
    
   
    
3. ₁₈Ar 
   
   See Answer
   Answer: 8
   
   There are two electrons in its 3s orbital and six electrons in its 3p orbitals.
   
    
   
    
4. ₁₉K 
   
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   Answer: 1
   
   There is one electron in its 4s orbital.
   
    
   
    
5. ₂₂Ti 
   
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   Answer: 2
   
   There are two electrons in its 4s orbital. The 3d electrons do not count as valence shell electrons.
   
    
   
    
6. ₂₈Ni 
   
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   Answer: 2
   
   There are two electrons in its 4s orbital. The 3d electrons do not count as valence shell electrons.
   
    
   
    
7. ₃₂Ge 
   
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   Answer: 4
   
   There are two electrons in its 4s orbital. There are two electrons in its 4p orbitals. (The 3d electrons do not count as valence electrons.)
   
    
   
    
8. ₃₄Se 
   
   See Answer
   Answer: 6
   
   There are two electrons in its 4s orbital. There are four electrons in its 4p orbitals. (The 3d electrons do not count as valence electrons.)
   
    
   
    
9. ₃₆Kr 
   
   See Answer
   Answer: 8
   
   There are two electrons in its 4s orbital. There are six electrons in its 4p orbitals. (The 3d electrons do not count as valence electrons.)
   
    
   
    

 
 
10. Is it possible for an electron to be a core electron in one atom and a valence shell electron in another atom?



 
11. A magnesium atom has 2 valence shell electrons. If it loses 2 electrons and becomes the Mg²⁺ ion, then how many valence shell electrons will it then have?