Valence Electrons. The electrons in the outermost shell are the valence electrons the electrons on an atom that can be gained or lost in a chemical reaction. Since filled d or f subshells are seldom disturbed in a chemical reaction, we can define valence electrons as follows: The electrons on an atom that are not present in the previous rare gas, ignoring filled d or f subshells.

In chemistry and physics, a valence electron is an outer shell electron that is associated with an atom, and that can participate in the formation of a chemical bond if the outer shell is not closed; in a single covalent bond, both atoms in the bond contribute one valence electron in order to form a shared pair. Total number of valence electrons in CN- = Valence electrons of Carbon + Valence electrons of Nitrogen + extra electron (as this molecule has a negative charge, it has accepted an extra electron which gives it a charge of -1 to this molecule) Carbon has four valence electrons. Nitrogen has five valence electrons. The number of valence electrons is the number of electrons in the outer shell, that the atom uses for bonding. Nitrogen has 5 electrons in its n=2 (outer) shell. There is a quick way of identifying the number of valence electrons - it is the same as the Group number (not for d-block elements, though). Nitrogen Number of Valence Electrons Nitrogen is an element that has 7 electrons and when we talk about the valence electrons then, valence electrons are those electrons that are present in the outer shell and are associated with molecules or an atom and which also has the capacity to participate in the chemical formation.

Lewis structure of NO₃^- ion is drawn step by step in this tutorial. Total valence electrons of nitrogen and oxygen atoms and negative charge are considered to draw the NO₃^- lewis structure. You will every fact of drawing lewis structures from this tutorial which will help you to draw more lewis structures in the future.

Lewis Structure of nitrite ion

Now, we are going to learn, how to draw this lewis structure of NO₃^- ion.

Steps of drawing NO₃^- lewis structure

Following steps are required to draw NO₃^- lewis structure and they are explained in detail in this tutorial.

1. Find total number of electrons of the valance shells of nitrogen and oxygen atoms and including charge of the anion
2. Total electrons pairs in valence shells
3. Center atom selection from nitrogen and oxygen atom
4. Put lone pairs on atoms
5. Stability of lewis structure - Check the stability and minimize charges on atoms by converting lone pairs to bonds.

Drawing correct lewis structure is important to draw resonance structures. In another tutorial, we learn how to draw resonance structures of nitrate ion.

Total number of electrons of the valance shells of nitrogen and oxygen atoms and charge of the anion

There are one nitrogen atom and three oxygen atoms in the nitrate ion. Also there is a -1 charge on the nitrate ion.

Nitrogen and oxygen are located at VA and VIA groups respectively in the periodic table. So nitrogen has five electrons in its valence shell. In oxygen atom, there are six electrons in its valence shell.

• Total valence electrons given by nitrogen atom = 5

There are three oxygen atoms in NO₃^-, Therefore

• Total valence electrons given by oxygen atoms = 6 *3 = 18

Due to -1 charge, another electrons is added

• Due to -1 charge, received electrons to valence electrons= 1

• Total valence electrons = 5 + 18 + 1 = 24

Total valence electrons pairs

Total valance electrons pairs = σ bonds + π bonds + lone pairs at valence shells

Total electron pairs are determined by dividing the number total valence electrons by two. For, NO₂^-, there are 24 valence electrons, so total pairs of electrons are 12.

Center atom of NO₂^-

To be the center atom, ability of having greater valance is important. Nitrogen can show valence,5. But, oxygen's maximum valence is 2. Therefore nitrogen has the more chance to be the center atom (See the figure). So, now we can build a sketch of NO₃^- ion.

Sketch of NO₂^- ion

Lone pairs on atoms

There are already three N-O bonds in the sketch. Therefore only nine valence electrons pairs are remaining to draw the rest of ion.

Start to mark those nine valence electrons pairs on outside atoms (oxygen atoms) as lone pairs. One oxygen atom will take three lone pairs following the octal rule (oxygen and nitrogen atoms cannot keep more than eight electrons in their valence shells). All nine valence electrons pairs (9) are spent when lone pairs are marked on oxygen atoms.

Therefore, there is no ine valence electrons pairs to mark on nitrogen atom.

Check the stability of drawn NO₂^- ion and minimize charges on atoms by converting lone pairs to bonds

Check charges on atoms and mark them as below. Charges are important to decide the lewis structure of the ion.

The drawn structure for NO₃^- is not a stable structure because oxygen atoms and nitrogen atoms have charges. When a molecule or ion has so many charges on atoms, that structure is not stable.

Now, we should try to minimize charges by converting lone pair or pairs which exist on oxygen atoms to bonds. So we convert one lone pair of one oxygen atom as a N-O bond.

Now there is a double bond between nitrogen and one oxygen atom. There are also two single bonds (N-O) with nitrogen atom and other oxygen atoms.

In new structure, charges of atoms are reduced. Now there is no any charge on one oxygen atom. Also, charge of nitrogen atom is reduced from +2 to +1. Now you understand this structure of NO₃^- is more stable than previous structure due to less charges on atoms.

But, We cannot convert more lone pairs of other oxygen atom to make a bond with nitrogen atom because nitrogen cannot keep more than eight electrons in its last valence shell.

Questions

How many lone pairs are around the nitrogen atom in nitrate ion?

No electrons pairs exist on nitrogen atom. But, on nitrogen atom, there is a +1 charge. Around the nitrogen atom, there are two single bonds and double bond.

Related tutorials

Quantum Numbers,
Atomic Orbitals, and
Electron Configurations

Contents:
Quantum Numbers and Atomic Orbitals
1. Principal Quantum Number (n)
2.Angular Momentum (Secondary, Azimunthal) Quantum Number (l)
3.Magnetic Quantum Number (m_l)
4.Spin Quantum Number (m_s)
Table of Allowed Quantum Numbers
Writing Electron Configurations
Properties of Monatomic Ions
References

Quantum Numbers and Atomic Orbitals

By solving the Schrödinger equation (Hy = Ey), we obtain a set of mathematical equations, called wave functions (y), which describe the probability of finding electrons at certain energy levels within an atom.

A wave function for an electron in an atom is called an atomic orbital; this atomic orbital describes a region of space in which there is a high probability of finding the electron. Energy changes within an atom are the result of an electron changing from a wave pattern with one energy to a wave pattern with a different energy (usually accompanied by the absorption or emission of a photon of light).

Each electron in an atom is described by four different quantum numbers. The first three (n, l, m_l) specify the particular orbital of interest, and the fourth (m_s) specifies how many electrons can occupy that orbital.

1. Principal Quantum Number (n): n = 1, 2, 3, …, ∞
   Specifies the energy of an electron and the size of the orbital (the distance from the nucleus of the peak in a radial probability distribution plot). All orbitals that have the same value of n are said to be in the same shell (level). For a hydrogen atom with n=1, the electron is in its ground state; if the electron is in the n=2 orbital, it is in an excited state. The total number of orbitals for a given n value is n².

2. Angular Momentum (Secondary, Azimunthal) Quantum Number (l): l = 0, ..., n-1.
   Specifies the shape of an orbital with a particular principal quantum number. The secondary quantum number divides the shells into smaller groups of orbitals called subshells (sublevels). Usually, a letter code is used to identify l to avoid confusion with n:

The subshell with n=2 and l=1 is the 2p subshell; if n=3 and l=0, it is the 3s subshell, and so on. The value of l also has a slight effect on the energy of the subshell; the energy of the subshell increases with l (s < p < d < f).

3. Magnetic Quantum Number (m_l): m_l = -l, ..., 0, ..., +l.
   Specifies the orientation in space of an orbital of a given energy (n) and shape (l). This number divides the subshell into individual orbitals which hold the electrons; there are 2l+1 orbitals in each subshell. Thus the s subshell has only one orbital, the p subshell has three orbitals, and so on.

4. Spin Quantum Number (m_s): m_s = +½ or -½.
   Specifies the orientation of the spin axis of an electron. An electron can spin in only one of two directions (sometimes called up and down).
   The Pauli exclusion principle (Wolfgang Pauli, Nobel Prize 1945) states that no two electrons in the same atom can have identical values for all four of their quantum numbers. What this means is that no more than two electrons can occupy the same orbital, and that two electrons in the same orbital must have opposite spins.
   Because an electron spins, it creates a magnetic field, which can be oriented in one of two directions. For two electrons in the same orbital, the spins must be opposite to each other; the spins are said to be paired. These substances are not attracted to magnets and are said to be diamagnetic. Atoms with more electrons that spin in one direction than another contain unpaired electrons. These substances are weakly attracted to magnets and are said to be paramagnetic.

Table of Allowed Quantum Numbers

Writing Electron Configurations

The distribution of electrons among the orbitals of an atom is called the electron configuration. The electrons are filled in according to a scheme known as the Aufbau principle ('building-up'), which corresponds (for the most part) to increasing energy of the subshells:

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

It is not necessary to memorize this listing, because the order in which the electrons are filled in can be read from the periodic table in the following fashion:

Or, to summarize:

In electron configurations, write in the orbitals that are occupied by electrons, followed by a superscript to indicate how many electrons are in the set of orbitals (e.g., H 1s¹)

Another way to indicate the placement of electrons is an orbital diagram, in which each orbital is represented by a square (or circle), and the electrons as arrows pointing up or down (indicating the electron spin). When electrons are placed in a set of orbitals of equal energy, they are spread out as much as possible to give as few paired electrons as possible (Hund's rule).

examples will be added at a later date

In a ground state configuration, all of the electrons are in as low an energy level as it is possible for them to be. When an electron absorbs energy, it occupies a higher energy orbital, and is said to be in an excited state.

Properties of Monatomic Ions

The electrons in the outermost shell (the ones with the highest value of n) are the most energetic, and are the ones which are exposed to other atoms. This shell is known as the valence shell. The inner, core electrons (inner shell) do not usually play a role in chemical bonding.

Elements with similar properties generally have similar outer shell configurations. For instance, we already know that the alkali metals (Group I) always form ions with a +1 charge; the 'extra' s¹ electron is the one that's lost:

The next shell down is now the outermost shell, which is now full — meaning there is very little tendency to gain or lose more electrons. The ion's electron configuration is the same as the nearest noble gas — the ion is said to be isoelectronic with the nearest noble gas. Atoms 'prefer' to have a filled outermost shell because this is more electronically stable.

• The Group IIA and IIIA metals also tend to lose all of their valence electrons to form cations.

Nitrogen Valence Electrons

• The Group IV and V metals can lose either the electrons from the p subshell, or from both the s and p subshells, thus attaining a pseudo-noble gas configuration.

• The Group IV - VII non-metals gain electrons until their valence shells are full (8 electrons).

• The Group VIII noble gases already possess a full outer shell, so they have no tendency to form ions.

• Transition metals (B-group) usually form +2 charges from losing the valence s electrons, but can also lose electrons from the highest d level to form other charges.

References

N=10 Valence Electrons

Martin S. Silberberg, Chemistry: The Molecular Nature of Matter and Change, 2nd ed. Boston: McGraw-Hill, 2000, p. 277-284, 293-307.

N Valence Electrons Number