How To Balance Chemical Equations - Steps & Example

Chemical reactions are happening all around us. They go in your body as you read this article, in all the plants and animals around you, and in almost every physio-chemical process. Different elements and compounds (the chemicals) react with each other to produce some new compound or chemical (the formation of new elements is nuclear, not chemical). Chemical reactions are all about atoms and molecules of chemicals reacting with other chemical atoms and molecules under specific conditions.

Chemical equations are written notations or descriptions of chemical reactions. Symbols of elements and compounds & the universal law of conservation of mass/matter are used to write chemical equations. Chemical equations give us a pretty good idea about the number of chemicals required to produce a new chemical through a reaction.

One key thing about chemical equations is that they MUST be balanced. But how do we do so? And why do we need to balance them? This article explains.

What does it mean to balance chemical equations?

Let’s start with the equation of a simple chemical reaction.

2H₂ + O₂ à 2H₂O

H₂ is a hydrogen molecule. The 2 as the subscript means that two hydrogen atoms combine to form a stable molecule. O₂ is an Oxygen molecule that’s also diatomic. Hydrogen and Oxygen, when ignited together, produce water under normal temperature and pressure. The chemical reaction also releases energy as the molecular bonds are broken.

Now, note the coefficient ‘2’ before the hydrogen atom. Well, placing it balances out the entire chemical equation. Here’s how:

• Multiply the coefficient 2 with the 2 in the subscript of the hydrogen atom. That gives you the total number of hydrogen molecules participating in the above chemical reaction.

2 * H₂ = 4 H

Two hydrogen molecules or 4 H atoms partake in the reaction.

• Just one Oxygen molecule or two O atoms are participating in the reaction.  
• Two molecules of H₂O are produced. If we multiply and add accordingly, we find that both sides of the expression balance out, turning it into an equation.

(2 * 2) H + 2 O = 2H₂O (2 * 2 H + 2 O)

or

2H₂ + O₂ à 2H₂O

This is the basic idea behind balancing chemical equations.

Chemical formulas and symbols on the left-hand side denote the starting materials in a chemical reaction. The formulas and symbols on the right-hand side denote the final products. The total number of atoms on the left MUST BE EQUAL to the total number on the right.

The coefficients and the symbols/formulas do not represent atoms and, also, not molecules– they are formula units. The units can be atoms, molecules, or moles. We will look into the idea behind formula units and moles further on.

Let’s first focus on why we need to balance chemical equations.

Why must we balance chemical equations?

The universe we live in has certain laws and constraints in place. These laws define everything – from the organic reactions on a remote exoplanet to the volcano experiment in your chemistry lab.

The law of conservation of mass and energy are two fundamental laws of our universe. According to them, matter and energy can neither be created nor destroyed! The amount of reactants (the starting materials) must produce an equivalent amount of products. That is the essence of the law of conservation of mass (and if we dig deeper, the law of conservation of energy).

The law of conservation of mass, a universal law, is the reason why we need to balance chemical equations.

In any chemical reaction, the chemical bonds between the atoms in a molecule break, and atoms rearrange to form new molecules. Atoms do not disappear, and no new atoms form. Hence, we must balance chemical equations and keep the number and type of atoms the same on both sides.

The coefficients before the symbols and formulas multiply the number of atoms. A balanced chemical equation is one where all the numbers and kinds of atoms on both sides of the equal sign are the same.

A Bit About The Law of Conservation of Mass & Energy

The law states that mass can never be created or destroyed.

Well, while this statement is not entirely true. Matter and anti-matter annihilate each other to produce energy. What’s universal is the conversion of mass into an equivalent quantum of energy. That is what Professor Albert Einstein’s infamous equation suggests à E = m * C².

An example where matter annihilates to produce energy is the combination of an electron and an anti-electron or positron.

Positron e⁺ + Electron e^- à g + g Energetic Gamma Rays

Other sub-atomic particles can also form under different conditions.

Matter and anti-matter are one another's counterparts. They share the same mass but are opposite in certain qualities, such as electric charge. Matter and anti-matter always appear in pairs, and destruction occurs when matter & anti-matter combine.

Annihilation or destruction of matter always occurs at the sub-atomic level. It is not exactly destruction but the transformation of matter into energy. And these transformations take place only at the sub-atomic level due to the relative instability of sub-atomic particles when compared to atoms.

Matter annihilation/transformation generally involves subatomic particles and occurs during nuclear reactions under extreme conditions. Regarding chemical reactions, the energies, conditions, and particles involved do not result in any matte annihilation. The law of conservation of mass/matter holds when chemical reactions occur. And therein lies the need to balance chemical reactions.

Let’s get back on track and take a closer look at the process of chemical equation balancing.  

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The Equation Balancing Process

There are four key steps of balancing equations à 

1. First, write down the unbalanced chemical equation. Use the right symbols and formulas for all the starting elements & compounds.
2. Use basic arithmetic skills and add appropriate coefficients to balance out the number of atoms of each element.
3. Check and re-check to ensure the numbers and kinds of atoms on both sides of the expression or equation are the same.
4. Always reduce the coefficients to their lowest value.

To balance everything, we must have the smallest ratio of whole numbers on both sides of the equation. We can use fractional coefficients in certain cases, though.

Here’s an example à 

Consider the following compound: C₃H₆O₂ (chemical name = methoxy acetaldehyde)

When combusted, the following reaction takes place:

C₃H₆O₂ + O₂ à CO₂ + H₂O

How will you balance the above?

• Initially, hydrogen and carbon get balanced if we put a 3 in front of the carbon dioxide and the water on the right-hand side.  

C₃H₆O₂ + O₂ à 3 CO₂ + 3 H₂O

But Oxygen, the great life-giver, loses its balance. There are 9 oxygen atoms on the right and 4 on the left.

• Everything balances out if we put 3.5 as a coefficient in front of the Oxygen molecule on the left-hand side.

C₃H₆O₂ +3.5 O₂ à 3 CO₂ + 3 H₂O

• We can also balance the above equation in the following way:

2 C₃H₆O₂ + 7 O₂ à 6 CO₂ + 6 H₂O

Only whole numbers and no fractions balance things here. However, the general rule is to use the lowest whole-number ratio. Use fractions only when no whole-number ratios can balance things out.

Now, you must wonder about the exact quantity of materials reacting and being produced. Is it just molecules? Is it grams or kilograms of the starting materials? Well, they are moles (not rodents).

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Moles & Stoichiometry

We know that the idea of balancing chemical equations is to have the same number of atoms on both sides. That’s why we must balance them using whole numbers or fractions. But when we are using fractions, does that mean only a fraction of the molecules of an element or compound is taking part in the reaction?

Not at all, as chemical formulas can depict several things, including moles, molecules, or atoms. That’s why, in chemistry, we use the term formula units.

Let’s go back to our first example.

2H₂ + O₂ à 2H₂O

The convention is to always use the lowest number ratio as balancing coefficients. But we can also balance things (and adhere to the law of conservation of mass) using other whole-number ratios.

4H₂ + 2O₂ à 4H₂O

or

22H₂ + 11O₂ à 22 H₂O

As can be seen, any multiple of the lowest number ratio (2:1:2) can balance the equation and preserve mass conversion.

Stoichiometry is the branch of chemistry that studies the numerical quantities, mathematical operations, and relationships between the reactants & products in a chemical reaction.

Stoichiometry defines Avogadro’s Number, a constant that acts as a normalization factor for defining the amount of substance in a chemical reaction. It equals 6.023 * 10²³, which is 1 mole (mol).

1 mole is a unit of measurement for the number of constituent particles in a system, which, in our case, is a chemical reaction.

1 mole can be equal to 6.023 * 10²³ atoms, molecules, ions, or any other particle, depending on the kind of reaction or system. Moles make it easy to understand how many particles participate in a chemical reaction. So, when balancing chemical equations, we balance the number of particles taking part.

Balanced chemical equations balance the number of moles or molar amounts of both reactants & products. For the equation,

2H₂ + O₂ à 2H₂O

When combined under the right set of conditions, the 2 moles of hydrogen and 1 mole of Oxygen produce 2 moles of water. As this is a molecular reaction, we consider a single mole representing 6.023 * 10²³ molecules.

However, it is best to keep it simple when it comes to balancing chemical equations. Always balance the number of atoms of every element on both sides and use the lowest whole-number ratio.

Now, it is time we looked at chemical reactions involving ions.

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Balancing Ionic Equations

Ions are electrically charged molecules or atoms with either an excess or a deficiency of electrons. Till now, we have looked only at molecular chemical reactions. The rules for balancing equations involving ions remain the same. Also, keep in mind that moles can represent both molecules or ions.

In actuality, all chemical reactions involve ions. Atoms or molecules lose and/or gain electrons to form new molecules, compounds, and products. Remember that chemical reactions cannot produce new atoms; nuclear reactions are necessary.

Suppose we are mixing two aqueous solutions of calcium chloride (CaCl₂) and silver nitrate (AgNO₃). The balanced equation of the molecular chemical reaction is:

CaCl₂ (aq) + 2 AgNO₃ (aq) à Ca (NO₃)₂ (aq) + 2AgCl (s)

The entire process involves ionic dissociation of the molecules. Here’s the breakdown:

CaCl₂ (aq) ⟶ Ca²⁺(aq) + 2Cl⁻(aq)

2AgNO₃ (aq) ⟶ 2Ag⁺(aq) + 2NO^{3 −}(aq)

Ca (NO3 )₂ (aq) ⟶ Ca²⁺(aq) + 2NO^{3 −}(aq)

The complete ionic equation will be something like this:

Ca²⁺(aq) + 2Cl⁻(aq) + 2Ag⁺(aq) + 2NO₃⁻(aq) ⟶ Ca²⁺(aq) + 2NO₃ ⁻(aq) + 2AgCl(s)

As can be seen, Ca²⁺ and 2NO₃⁻ remain the same. These are spectator ions that aid in maintaining charge neutrality in the entire system. Ag and Cl combine to produce 2 moles of silver chloride.

Next, it is time to look at some of the most common types of chemical equations in chemistry & stoichiometry, acid-base reactions.

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Balancing Acid-Base Chemical Reactions

Simply put, acids are chemical compounds that release positive hydrogen ions (H⁺), and bases are compounds that produce negative hydroxide ions (OH^-). Hydrogen ions are too active to exist in ionic form and generally combine with H₂O to form hydronium ions (H₃O⁺).

Acids and bases are consumed by us in different shapes & forms every day. And acid-base reactions occur all around us. In every case, a hydrogen ion is transferred from one compound or chemical to the other. The chemical reactions between acids and bases lead to the production of hydronium and hydroxide ions. The presence of these ions is used to classify any chemical reaction as an acid-base reaction. Acids and bases neutralize each other and produce salt & water in a neutralization process.

Here’s an example:

Mg (OH)₂ + 2HCl à  MgCl₂ + 2H₂O

This is a reaction that takes place in our stomach when we ingest antacids to counter acidity. Mg (OH)₂, known as milk of magnesia, is the key compound in all antacids. The human stomach produces hydrochloric acid HCl to help digest food. When binging on oily, fried, and heavy foods causes acidity, the milk of magnesia in antacids neutralizes the acid in the stomach. The resultant products are salt, MgCl_2, and water.

Balancing chemical equations needs practice. Balance chemical equations as often as possible and of different types, including oxidation-reduction reactions.

Balancing Oxidation-Reduction Reactions

Reduction-oxidation reactions or Redox reactions involve the addition of an oxidizing agent to one chemical compound and removal from another. Oxygen is not the only oxidization agent but any chemical agent that removes electrons or negatively charged ions from a compound. Reduction reactions do the opposite, removing oxygen atoms or adding electrons or positively charged ions.

Here’s an example:

2Na(s) + Cl₂ (g) ⟶ 2NaCl(s)

2Na(s) ⟶ 2Na⁺(s) + 2e⁻ 

Cl₂ (g) + 2e⁻ ⟶ 2Cl⁻(s)

The sodium ion loses electrons and, thus, gets oxidized. Chlorine gains electrons and gets reduced. This loss and gain of electrons create the complementary nature of redox reactions. Sodium is the reducing agent, while chlorine is the oxidizing agent. Balancing redox reactions is the same as balancing any other chemical reaction.

The above reaction between sodium and chloride is ionic or electrovalent, hence, the explicit transfer of electrons. In the case of covalent reactions, redox processes may not involve any transfer of electrons but the addition/removal of Oxygen.  

Several reactions can be classified as redox reactions, such as combustion reactions. Fuels act as reductants or reducing agents and react vigorously with oxidants to produce heat, light, & immense thrust. The combustion of solid rocket fuels is a classic example of powerful redox reactions.

Ever wondered what propellant and chemical reaction produces such immense thrust in a rocket? It is a major redox reaction involving solid rocket propellants, liquid hydrogen fuel, and liquid oxygen oxidizers.

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Equations of Two Other Major Chemical Reaction Types

Displacement Reactions

Dip a clean iron nail in a solution of copper sulfate for some time. You will notice that the nail becomes coppery while the blue of the copper sulfate fades. Iron has displaced copper from the sulfate solution to form iron sulfate. The displaced copper gets deposited on the nail.

The equation for the chemical reaction is:

Fe + CuSO₄ à FeSO₄ + Cu

Double Displacement Reaction

A white precipitate will form if you mix sodium sulfate with barium chloride. The reaction is a double displacement reaction where the following takes place:

Na₂SO₄ + BaCl₂ à BaSO₄ + 2NaCl

Barium displaces sodium from sodium sulfate while itself getting displaced by barium chloride. The white precipitate formed is Barium Sulphate. We have looked into the preceding sections' redox and decomposition reactions.

Combination Reactions

Combination reactions involve reactants combining to produce a new product. Reactive elements such as Oxygen easily combine with other elements. A very common example of a combination reaction is the rusting of iron.

The chemical equation for the rusting reaction is:

4 Fe + 3 O₂ à 2 Fe₂O₃ 

Ferrous oxide is the product formed and is an ionic compound since it’s formed from metal & non-metal. Any metal can combine with Oxygen to form its oxide. The more reactive the metal (Potassium, Sodium, Magnesium), the more energetic the reaction. And not just metals & non-metals, but non-metals combine to form molecular compounds.

For example, Sulphur burns to form Sulphur dioxide as per the following chemical reaction:

2S + 3O₂ à 2SO₃

We wrap this write-up with a quick review of the generic chemical equation balancing process.

Interpreting Chemical Reactions Through Equations

Balancing a chemical equation can reveal important information about a reaction's relative quantity of materials. The coefficient of every reactant and product is the number of formula units participating in the reaction. It can be moles, molecules, ions, etc. Let’s look at an example.

(NH₄)₂Cr₂O₇ à Cr₂O ₃ + N₂ + 4H₂O

• If we look at the ionic reaction, two NH₄⁺ions and one Cr₂O₇^- ions yield 1 formula unit of Cr₂O ₃, 1 formula unit of nitrogen, and 4 formula units of water molecules.
• One mole of (NH₄)₂Cr₂O₇ produces 1 mole of Cr₂O₃, 1 mole of nitrogen, and four moles of water.
• Again, if we consider Avogadro’s number, then 6.023 * 10²³formula units (moles, molecules, or ions) produces 6.023 * 10²³ formula units of Cr₂O ₃, 6.023 * 10²³ formula units of nitrogen, and 4 * 6.023 * 10²³ units of water.

All of the above statements are correct and completely chemically equivalent. The ratio between the number of moles of one reactant with another or with a product or amongst two products is known as the mole ratio.

The quantity of reactants and products can be easily interpreted from a balanced chemical equation. However, you cannot find out anything about the reaction rate or the kind and amount of energy released during the reaction.

Let’s wrap things up with some handy tips on balancing chemical equations.

Some Extra Chemical Equations Balancing Pointers

Many advanced organic and inorganic reactions can involve gigantic formulas and equations. Many people struggle with balancing such large equations. That’s where online balancing chemical equations with online solvers come in handy.

And that finally brings us to the end of this write-up. Hope this article is a handy guide for anyone struggling with formulating and balancing chemical equations. Practice often, use equation solvers, and if you need some help with your chemistry assignments, connect with the experts at AllEssayWriter.com.