Valence Electrons

Valence Electrons Definition and Basics

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.

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.

When we look at an element's oxidation state, it tells us a key fact about that element's electrons.

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.

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.

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).

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.

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.

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:

In summary, the valence electrons for a neutral atom are:

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,

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.

Let's do an example:

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.