N Valence Electrons

1. Valence Electrons Rings
2. List Of Valence Electrons For Each Element
3. N=10 Valence Electrons

You can find valence electrons with a shortcut using the periodic table, but it's good to only do that after you understand why the shortcut works, and to do. 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, NO 2-, there are 18 valence electrons pairs, so total pairs of electrons are 9.

Nitrogen is present in almost all proteins and plays important roles in both biochemical applications and industrial applications. Nitrogen forms strong bonds because of its ability to form a triple bond with its self, and other elements. Thus, there is a lot of energy in the compounds of nitrogen. Before 100 years ago, little was known about nitrogen. Now, nitrogen is commonly used to preserve food, and as a fertilizer.

Introduction

Nitrogen is found to have either 3 or 5 valence electrons and lies at the top of Group 15 on the periodic table. It can have either 3 or 5 valence electrons because it can bond in the outer 2p and 2s orbitals. Molecular nitrogen ((N_2)) is not reactive at standard temperature and pressure and is a colorless and odorless gas.

Nitrogen is a non-metal element that occurs most abundantly in the atmosphere, nitrogen gas (N₂) comprises 78.1% of the volume of the Earth's air. It only appears in 0.002% of the earth's crust by mass. Compounds of nitrogen are found in foods, explosives, poisons, and fertilizers. Nitrogen makes up DNA in the form of nitrogenous bases as well as in neurotransmitters. It is one of the largest industrial gases, and is produced commercially as a gas and a liquid.

History

Nitrogen, which makes up about 78% of our atmosphere, is a colorless, odorless, tasteless and chemically unreactive gas at room temperature. It is named from the Greek nitron + genes for soda forming. For many years during the 1500's and 1600's scientists hinted that there was another gas in the atmosphere besides carbon dioxide and oxygen. It was not until the 1700's that scientists could prove there was in fact another gas that took up mass in the atmosphere of the Earth.

Discovered in 1772 by Daniel Rutherford (and independently by others such as Priestly and Cavendish) who was able to remove oxygen and carbon dioxide from a contained tube full of air. He showed that there was residual gas that did not support combustion like oxygen or carbon dioxide. While his experiment was the one that proved that nitrogen existed, other experiments were also going in London where they called the substance 'burnt' or 'dephlogisticated air'.

Nitrogen is the fourth most abundant element in humans and it is more abundant in the known universe than carbon or silicon. Most commercially produced nitrogen gas is recovered from liquefied air. Of that amount, the majority is used to manufacture ammonia ((NH_3)) via the Haber process. Much is also converted to nitric acid ((HNO_3)).

Isotopes

Nitrogen has two naturally occurring isotopes, nitrogen-14 and nitrogen-15, which can be separated with chemical exchanges or thermal diffusion. Nitrogen also has isotopes with 12, 13, 16, 17 masses, but they are radioactive.

• Nitrogen 14 is the most abundant form of nitrogen and makes up more than 99% of all nitrogen found on Earth. It is a stable compound and is non-radioactive. Nitrogen-14 has the most practical uses, and is found in agricultural practices, food preservation, biochemicals, and biomedical research. Nitrogen-14 is found in abundance in the atmosphere and among many living organisms. It has 5 valence electrons and is not a good electrical conductor.
• Nitrogen-15 is the other stable form of nitrogen. It is often used in medical research and preservation. The element is non-radioactive and therefore can also be sometimes used in agricultural practices. Nitrogen-15 is also used in brain research, specifically nuclear magnetic resonance spectroscopy (NMR), because unlike nitrogen-14 (nuclear spin of 1), it has a nuclear spin of 1/2 which has benefits when it comes to observing MRI research and NMR observations. Lastly, nitrogen-15 can be used as label or in some proteins in biology. Scientists mainly use this compound for research purposes and have not yet seen its full potential for uses in brain research.

Compounds

The two most common compounds of nitrogen are Potassium Nitrate (KNO₃) and Sodium Nitrate (NaNO₃). These two compounds are formed by decomposing organic matter that has potassium or sodium present and are often found in fertilizers and byproducts of industrial waste. Most nitrogen compounds have a positive Gibbs free energy (i.e., reactions are not spontaneous).

The dinitrogen molecule ((N_2)) is an 'unusually stable' compound, particularly because nitrogen forms a triple bond with itself. This triple bond is difficult hard to break. For dinitrogen to follow the octet rule, it must have a triple bond. Nitrogen has a total of 5 valence electrons, so doubling that, we would have a total of 10 valence electrons with two nitrogen atoms. The octet requires an atom to have 8 total electrons in order to have a full valence shell, therefore it needs to have a triple bond. The compound is also very inert, since it has a triple bond. Triple bonds are very hard to break, so they keep their full valence shell instead of reacting with other compounds or atoms. Think of it this way, each triple bond is like a rubber band, with three rubber bands, the nitrogen atoms are very attracted to each other.

Nitrides

Nitrides are compounds of nitrogen with a less electronegative atom; in other words it's a compound with atoms that have a less full valence shell. These compounds form with lithium and Group 2 metals. Nitrides usually has an oxidation state of -3.

[3Mg + N_2 rightarrow Mg_3N_2 label{1}]

When mixed with water, nitrogen will form ammonia and, this nitride ion acts as a very strong base.

[N + 3H_2O_{(l)} rightarrow NH_3 + 3OH^-_{(aq)} label{2}]

When nitrogen forms with other compounds it primarily forms covalent bonds. These are normally done with other metals and look like: MN, M₃N, and M₄N. These compounds are typically hard, inert, and have high melting points because nitrogen's ability to form triple covalent bonds.

Ammonium Ions

Nitrogen goes through fixation by reaction with hydrogen gas over a catalyst. This process is used to produce ammonia. As mentioned earlier, this process allows us to use nitrogen as a fertilizer because it breaks down the strong triple bond held by N_2. The famous Haber-Bosch process for synthesis of ammonia looks like this:

[N_2 + 3H_2 rightarrow 2NH_3 label{3}]

Ammonia is a base and is also used in typical acid-base reactions.

Sophos xg 125. [2NH_{3(aq)} + H_2SO_4 rightarrow (NH_4)_2SO_{4(aq)} label{4}]

Nitride ions are very strong bases, especially in aqueous solutions.

Oxides of Nitrogen

Nitrides use a variety of different oxidation numbers from +1 to +5 to for oxide compounds. Almost all the oxides that form are gasses, and exist at 25 degrees Celsius. Oxides of nitrogen are acidic and easily attach protons.

[N_2O_5 + H_2O rightarrow 2HNO_{3 (aq)} label{5}]

The oxides play a large role in living organisms. They can be useful, yet dangerous.

• Dinitrogen monoxide (N₂O) is a anesthetic used at the dentist as a laughing gas.
• Nitrogen dioxide (NO₂) is harmful. It binds to hemoglobin molecules not allowing the molecule to release oxygen throughout the body. It is released from cars and is very harmful.
• Nitrate (NO₃^-) is a polyatomic ion.
• The more unstable nitrogen oxides allow for space travel.

Hydrides

Hydrides of nitrogen include ammonia (NH₃) and hyrdrazine (N₂H₄).

• In aqueous solution, ammonia forms the ammonium ion which we described above and it has special amphiprotic properties.
• Hyrdrazine is commonly used as rocket fuel

Applications of Nitrogen

• Nitrogen provides a blanketing for our atmosphere for the production of chemicals and electronic compartments.
• Nitrogen is used as fertilizer in agriculture to promote growth.
• Pressurized gas for oil.
• Refrigerant (such as freezing food fast)
• Explosives.
• Metals treatment/protectant via exposure to nitrogen instead of oxygen

References

1. Petrucci, Ralph H, William Harwood, and F. Herring. General Chemistry: Principles and Modern Applications. 8^th Ed. New Jersey: Pearson Education Inc, 2001.
2. Sadava, David et al. LIFE: The Science of Biology. Eighth Edition. Sinauer Associate.
3. Thomas, Jacob. Nitrogen and its Applications to Modern Future. San Diego State University Press: 2007.

Problems

• Complete and balance the following equations

N₂+ ___H₂→ ___NH_{_}

H₂N₂O₂ → ?

2NH₃ + CO₂ → ?

__Mg + N₂ → Mg_{_}N_{_}

N₂H₅ + H₂O → ?

• What are the different isotopes of Nitrogen?
• List the oxiadation states of various nitrogen oxides: N₂O, NO, N₂O_3, N₂O_4, N₂O₅
• List the different elements that Nitrogen will react with to make it basic or acidic..
• Uses of nitrogen

Answers

• Complete and balance the following equations

N₂+ 3H₂→ 2NH₃(Haber process)

H₂N₂O₂ → HNO

2NH₃ + CO₂ → (NH₂)₂CO + H₂O

2Mg + 3N₂ → Mg₃N₂

N₂H₅ + H₂O → N₂+ H⁺ + H₂O

• What are the different isotopes of Nitrogen?

Stable forms include nitrogen-14 and nitrogen-15

• List the oxidation states of various nitrogen oxides: +1, +2, +3, +4, +5 respectively

• List the different elements that Nitrogen will react with to make it basic or acidic :Nitride ion is a strong base when reacted with water, Ammonia is generally a weak acid
• Uses of nitrogen include anesthetic, Refrigerant, metal protector

Lewis structure of NO₂^- ion is drawn in this tutorial. Total valence electrons of nitrogen and oxygen atoms and negative charge also should be considered in the drawing of NO₂^- lewis structure.

• In aqueous solution, ammonia forms the ammonium ion which we described above and it has special amphiprotic properties.
• Hyrdrazine is commonly used as rocket fuel

Applications of Nitrogen

• Nitrogen provides a blanketing for our atmosphere for the production of chemicals and electronic compartments.
• Nitrogen is used as fertilizer in agriculture to promote growth.
• Pressurized gas for oil.
• Refrigerant (such as freezing food fast)
• Explosives.
• Metals treatment/protectant via exposure to nitrogen instead of oxygen

References

1. Petrucci, Ralph H, William Harwood, and F. Herring. General Chemistry: Principles and Modern Applications. 8^th Ed. New Jersey: Pearson Education Inc, 2001.
2. Sadava, David et al. LIFE: The Science of Biology. Eighth Edition. Sinauer Associate.
3. Thomas, Jacob. Nitrogen and its Applications to Modern Future. San Diego State University Press: 2007.

Problems

• Complete and balance the following equations

N₂+ ___H₂→ ___NH_{_}

H₂N₂O₂ → ?

2NH₃ + CO₂ → ?

__Mg + N₂ → Mg_{_}N_{_}

N₂H₅ + H₂O → ?

• What are the different isotopes of Nitrogen?
• List the oxiadation states of various nitrogen oxides: N₂O, NO, N₂O_3, N₂O_4, N₂O₅
• List the different elements that Nitrogen will react with to make it basic or acidic..
• Uses of nitrogen

Answers

• Complete and balance the following equations

N₂+ 3H₂→ 2NH₃(Haber process)

H₂N₂O₂ → HNO

2NH₃ + CO₂ → (NH₂)₂CO + H₂O

2Mg + 3N₂ → Mg₃N₂

N₂H₅ + H₂O → N₂+ H⁺ + H₂O

• What are the different isotopes of Nitrogen?

Stable forms include nitrogen-14 and nitrogen-15

• List the oxidation states of various nitrogen oxides: +1, +2, +3, +4, +5 respectively

• List the different elements that Nitrogen will react with to make it basic or acidic :Nitride ion is a strong base when reacted with water, Ammonia is generally a weak acid
• Uses of nitrogen include anesthetic, Refrigerant, metal protector

Lewis structure of NO₂^- ion is drawn in this tutorial. Total valence electrons of nitrogen and oxygen atoms and negative charge also should be considered in the drawing of NO₂^- lewis structure.

Now, we are going to learn, how to draw this lewis structure.

Steps of drawing NO₂^- lewis structure

Following steps are required to draw NO₂^- 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 charge of the anion
2. Total electrons pairs
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.

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

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

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 two oxygen atoms in NO₂, Therefore

• Total valence electrons given by oxygen atoms = 6 *2 = 12

Due to -1 charge, another electrons is added

• Due to -1 charge, received electrons = 1

• Total valence electrons = 5 + 12 + 1 = 18

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 18 valence electrons pairs, so total pairs of electrons are 9.

Center atom of NO₂^-

To be the center atom, ability of having greater valance is important. Therefore nitrogen has the more chance to be the center atom (See the figure). So, now we can build a sketch of NO₂^- ion.

Lone pairs on atoms

There are already two N-O bonds in the sketch. Therefore only seven valence electrons pairs are remaining.

Start to mark those seven 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).

Two oxygen atoms will take six valence electrons pairs. Now one valence electrons pair is remaining. Mark that remaining one on nitrogen atom.

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

The drawn structure for NO₂^- is not a stable one because both oxygen atoms and nitrogen atoms have charges.

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

Now there is a double bond between nitrogen and one oxygen atom. There is a single bond also with nitrogen atom and other oxygen atom.

In new structure, charges of atoms reduced. Now there is no any charge on one oxygen atom and nitrogen atom. Now you understand this structure of NO₂^- is more stable than previous structure due to less charges.

Valence Electrons Rings

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.

List Of Valence Electrons For Each Element

N=10 Valence Electrons

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