Electrons are some pretty important particles. They flow through our wires and cables to give us electricity, and they are also responsible for bonding! But did you know that there are actually two "classes" of electrons? In this article, we will be talking about valence electrons. We will learn what they are, how to count them, and learn why they are so important.
- This article is about valence electrons.
- First, we will learn about the basics of valence electrons.
- Then we will walk through the number of valence electrons for different elements/groups, starting with carbon.
- Next, we will learn how to use the periodic table to count valence electrons.
- Lastly, we will learn about orbitals and why they are important for understanding the valence electrons of transition metals.
Valence Electrons Definition and Basics
Valence electrons are the electrons that exist in the outermost/ highest energy level of an atom. These are the electrons that participate in bonding.
Here is a diagram of what valence electrons look like in the atom boron:
Fig. 1 - Boron's valence electrons
In our figure, the valence electrons are in purple, and the core electrons (any electron not in the highest energy level) are in yellow. For many atoms, the first energy level/shell can hold 2 electrons, then all other levels can hold 8.
This is true for all p-block atoms. When looking at d-block or f-block metals, things start to change, and there can be more than 8 valence electrons. For now, we will just look at atoms which prefer to hold up to 8 valence electrons.
Atoms want to be as stable as possible, so they want to have a full shell of valence electrons. We call this an octet since there are 8 electrons. Hydrogen and helium are two exceptions to the octet rule, since they only need two electrons to fill their valence electron shell. This is because they only have one energy level.
The noble gases are the most stable elements since they have full valence electron shells. In a sense, all atoms are trying to be like noble gases
When we look at an element's oxidation state, it tells us a key fact about that element's electrons.
Oxidation state is the theoretical charge an element will have if all of its bonds are fully ionic (electrons are either donated or taken, not shared).
The oxidation state will tell us:
- +x charge --> lose x electrons.
- -x charge --> gain x electrons.
So, an atom with an oxidation state of +1 will be missing 1 electron.
Elements usually have 1 or 2 oxidation states that are common due to the number of valence electrons it has in its neutral state.
For example, boron has 3 valence electrons. This means it has two options: lose three electrons to have a full shell, or gain 5 electrons to have a full shell. Since atoms will always take the "easier" route, boron will lose three electrons, which is why it is common to have a +3 oxidation state.
As another example, hydrogen can have a +1 or -1 oxidation state. Since it only has 1 valence electron, it can choose to fill or lose its shell. The +1 state is more common, but there are compounds that contain hydrogen in a state of -1 (such as BeH2, where Be has a +2 oxidation state). Hydrogen is known as a proton when it is +1, and a hydride when it is -1.
From here on I will be referring to the oxidation state as "charge", since it is the theoretical charge the element has.
Charged species
Typically speaking, elements will bond with other species so they can fill their octet. However, this is not always the case.
Boron has 3 valence electrons, so it can form 3 bonds and donate those three electrons. However, that doesn't always mean that boron will form three bonds. For example, boron can form 4 bonds with hydrogen to make BH4+.
Carbon valence electrons
Carbon is the 6th element in the periodic table. All neutral elements will have the same number of protons and electrons, which is equal to the atomic number. Since carbon has 6 electrons in total, it will have 4 valence electrons as shown below:
Fig. 2 - Carbon has 4 valence electrons
Because of its charge, carbon can form 4 bonds. Since there are two electrons held per bond, this means that with each bond it either gains an electron or loses an electron. Unless the two atoms are the same, the electrons will not be held equally between the two.
Thus, with every bond that is not a C–C bond, carbon will either be gaining or losing electrons. This equates to carbon having an oxidation state somewhere between -4 and +4, depending on what it's bonded to.
There are many examples of carbon having different oxidation states. Carbon dioxide, CO2, shows carbon with an oxidation state of +4. Carbon dioxide is a very common gas in the atmosphere which is linked to the greenhouse effect.
Methane, CH4, is a molecule that has carbon with an oxidation state of -4. Interestingly, methane is also a greenhouse gas. These two carbon atoms are in vastly different electronic environments, yet they both behave similarly.
All elements in group 14 (carbon group) have 4 valence electrons, and will try to form 4 bonds. However, as you move down the group, the properties of each element begin to change drastically.
Nitrogen valence electrons
Now onto nitrogen. Nitrogen has an atomic number of 7 (being in group 5), so it has 5 valence electrons as shown below:
Fig. 3 - Nitrogen has 5 valence electrons
Since nitrogen has 5 valence electrons, it can have a -3 charge. Nitrogen wants 3 more electrons to fill its octet, since that is easier than losing 5.
Since it lacks 3 electrons, nitrogen will generally form 3 bonds. Since it wants to gain electrons, it forms covalent bonds (6 electrons shared in total). The "extra" two electrons form a lone pair (a pair of electrons which are typically not shared with another atom).
Electrons form lone pairs because they occupy orbitals which are lower in energy. These are typically called non-bonding electrons. Although they can bond with other atoms, they will be the last of the valence electrons to do so.
Elements in nitrogen's group (group 15) have 5 valence electrons and will preferentially form 3 bonds.
Oxygen Valence Electrons
Oxygen's atomic number is 8 (group 6), so it has 6 valence electrons as shown below:
Fig. 4 - Oxygen has 6 valence electrons
Because it has 6 valence electrons, oxygen will generally have a -2 charge. This also means it can form two covalent bonds (4 total electrons shared), and the extra 4 electrons form two lone pairs.
Sulfur Valence Electrons
Sulfur is in the same group as oxygen, so it also has 6 valence electrons. The biggest difference between sulfur and oxygen is that sulfur has one more electron shell, as shown below:
Fig. 5 - Sulfur has 6 valence electrons, but 2 shells of core electrons.
Any electrons which are not in the outermost valence shell are known as core electrons. Larger atoms will have more core electrons to balance the larger positive charge from a larger nucleus.
Even though sulfur has another level of core electrons, that doesn't affect its number of valence electrons. Sulfur may also have a charge of -2, which will form 2 covalent bonds, and can have two lone pairs. However, unlike oxygen, sulfur may possess much higher oxidation states (more positive).
Although oxygen typically only forms 2 bonds, sulfur is not so strict. For example, sulfur hexafluoride (SF6), is known as a hypervalent molecule. Although it is not extremely common, this is sulfur in a +6 oxidation state.
As mentioned previously, all elements within a group will have the same number of valence electrons and will have similar bonding properties. Even though going down a group adds a shell of electrons, this doesn't affect the number of valence electrons.
While core electrons don't mean much when it comes to bonding, they do affect other factors due to the shielding effect. The farther away valence electrons are from the nucleus, the less "pull" they feel from it. This is called the shielding effect. The more core electrons, the stronger this effect is. This affects properties such as ionization energy, which is the energy it takes to move one valence electron. The farther you go down a group (i.e. the more core electrons), the easier it is to remove an electron.
Valence Electrons Periodic Table
We can figure out the number of valence electrons an element has just by looking at the periodic table.
Fig. 6 - The periodic table.
The number of valence electrons is equal to the group number (ignoring transition metals). The main exception to this is helium, which only has 2 valence electrons.
Let's work on a problem:
How many valence electrons are in the following atoms: a) Cl b) Ga c) Sr
For Chlorine, it is in the 7th group across, so it has seven valence electrons.
For Gallium, it is the 3rd group across, so it has three valence electrons.
For Strontium, it is in the 2nd group across, so it has two valence electrons.
In summary, the valence electrons for a neutral atom are:
| Group Number | Number of Valence electrons |
| Group 1 | 1 |
| Group 2 | 2 |
| Group 13 | 3 |
| Group 14 | 4 |
| Group 15 | 5 |
| Group 16 | 6 |
| Group 17 | 7 |
| Group 18 | 8 |
This method only applies to neutral atoms. If an atom has a charge, the formula is:
$$\text{number of valence electrons}=\text{number of valence electrons in neutral atom}-\text{charge}$$
Let's work on some examples,
how many valence electrons do the following ions have? a) O2- b) Mg+
the valence electrons for O2- is:
$$\text{number of valence electrons}=6-(-2)$$
$$\text{number of valence electrons}=8$$
And for Mg+:
$$\text{number of valence electrons}=2-1$$
$$\text{number of valence electrons}=1$$
Valence Electrons and Orbitals
Up until now, I have been referring to electrons being housed in "shells". This is a simplification of what is actually happening. Electrons are in orbitals. Orbitals are regions that electrons may be orbiting in. The main orbital types from least to greatest energy are:
- S-orbitals
- Has 1 subshell that holds 2 electrons
- P-orbitals
- Has 3 subshells, each holding 2 electrons, for a total of 6 electrons
- D-orbitals
- Has 5 subshells, each holding 2 electrons, for a total of 10 electrons
- F-orbitals
- Has 7 subshells, each holding 2 electrons, for a total of 14 electrons
Below is the periodic table with the labeled orbitals
Fig. 7 - Periodic table with orbitals
Each period is its own energy level. P-orbitals start appearing in period 2, and d-orbitals start appearing in period 4 (though they start counting at 3). F-orbitals start appearing in the lanthanides and actinides (the separated two rows).
The way we count our electrons is by moving from right to left, starting at the beginning of the table. For example, carbon has an electron configuration of 1s22s22p2.
So what does this have to do with our shells? The first "shell" represents the 1s orbital. The shells after (ignoring transition metals) represent the s and p-orbitals. For example, oxygen has 6 valence electrons, these six electrons fill up the 2s orbital, and partially fill the 2p subshells (2s22p4).
Looking at the orbitals explains how valence electrons work for transition metals. For non-transition metals, we count to 8, but for transition metals, we count to 12. The valence electrons for transition metals are equal to the number of s-electrons plus the number of d-electrons. In other words, the number of valence electrons for a transition metal is equal to how many spaces across the periodic table it is.
18-electron rule
For metal complexes, transition metals can hold 18 electrons. Like with the other elements, transition metals wants to have a full set of valence electrons, like the octet rule with noble gases. For elements in period 4 and below, they can have valence electrons in the s-, p-, and d-orbitals, for a total of 18 valence electrons.
A very common transition metal catalyst is Pd(PPh3)4. Palladium (Pd) is in group 10, so it has 10 valence electrons. It receives 2 electrons per triphenylphosphine (PPh3). Therefore, the entire complex will have 18 valence electrons.
Let's do an example:
How many valence electrons does nickel (Ni) have?
Nickel is 10 across on the periodic table, so it has 10 valence electrons. This also means it has 8 3d-electrons (is in the 8th transition metal group) and 2 4s-electrons.
Essentially, valence electrons are whatever electrons exist in the highest orbitals. P-orbitals are higher in energy than d-orbitals, which is why the d-electrons aren't valence electrons for non-transition metals.
Valence Electrons - Key takeaways
- Valence electrons are the electrons that exist in the outermost/the highest energy level of an atom. These are the electrons that participate in bonding.
- The charge on an ion tells us how many valence electrons that atom wants to lose or gain.
- +x charge --> lose x electrons.
- -x charge --> gain x electrons.
- When electrons are donated from one atom to another through a bond, that is called an ionic bond. When the electrons within a bond are shared between the two atoms, that is called a covalent bond.
- For neutral atoms, the number of valence electrons is equal to the group number (excluding valence electrons).
- For charged atoms, the # is equal to the neutral number minus the charge.
- For transition metals, the number of valence electrons is equal to the group number.
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