Chemical Reactions Worksheets (Explanation and Examples)

Chemical reactions are occurring all around us. Some everyday examples are –

1. The rusting of iron, where iron combines with oxygen to form a red-colored oxide
2. Combustion process when you strike a match
3. Photosynthesis process in plants wherein carbon dioxide and water convert into glucose and oxygen
4. When vinegar and baking soda combine to form a bubbling mixture
5. Batteries which convert stored chemical energy into electrical energy

Chemical Equations

Chemical equations are a shorthand representation of what occurs in chemical reactions. Equations are written with chemical symbols and numbers to explain the reactants and products involved in a reaction. The reactants are shown on the left side of the equation while the products are shown on the right. An arrow in the middle means “yields” or “produces”.

For example –

2H₂(g)+ O₂(g)→ 2H₂O(l)

This equation explains that two molecules of hydrogen gas(g)react with one molecule of oxygen gas(g) to yield two molecules of a compound called water (represented as H₂O), which is liquid (l).

A chemical reaction between ammonia gas (NH₃) and oxygen gas which yields Nitric Acid (HNO₃) and water is represented as –

NH₃(g)+ 2O₂(g)→ HNO₃(aq)+ H₂O(l)

This equation explains that one molecule of ammonia gas(g)reacts with two molecules of oxygen gas(g)to yield one molecule of nitric acid, which is an aqueous solution (aq), and one molecule of water, which is liquid (l). (An aqueous solution simply means that it is ‘dissolved in the water’).

Balancing Chemical Equations

Chemical equations need to be balanced. This follows the law of conservation of matter wherein matter can neither be created nor destroyed as a result of a chemical reaction. In other words, the number of atoms of each element in an equation must remain the same on either side of that equation.

In the equation to create water, if we simply wrote:

H₂(g)+ O₂(g)→ H₂O (l) and counted the atoms on either side, we’d get:

Left side: 2 atoms of hydrogen (represented as H₂) and 2 atoms of oxygen (represented as O₂)

Right side: 2 atoms of hydrogen (represented as H₂) but only one atom of oxygen (O).

This is an unbalanced equation!

To correctly balance equations we make use of Coefficients. A number written before an element is known as a coefficient.

Since oxygen atoms are unbalanced, we add “2” before the water molecule H₂O to correctly show 2 atoms of oxygen on the right.

(Note 1: An equation cannot be balanced by adding or removing a subscript – for example, adding a subscript to oxygen on the right side to show “H₂O₂” might balance the equation but it is chemically incorrect. This is because H₂O₂ is a completely different substance – hydrogen peroxide – compared to water – H₂O).

(Note 2: The coefficient “2” before H₂O applies to both H₂ and to O. The coefficient is multiplied by the number of atoms in the subscript of the element to give us the total number of atoms for that element. This coefficient therefore gives us 4 atoms of hydrogen and 2 atoms of oxygen on the right side of the equation).

Thus:H₂(g)+ O₂(g)→ 2H₂O (l) helps balance the oxygen atoms on either side but leaves the hydrogen atoms unbalanced. The hydrogen count of atoms is 2 on the left side and 4 on the right. The equation is still unbalanced!

To completely balance the equation, we need yet another coefficient “2” on the left side before the hydrogen element.

Thus: 2H₂(g)+ O₂(g)→ 2H₂O (l) which now correctly gives –

Left side: 4 atoms of hydrogen and 2 atoms of oxygen

Right side: 4 atoms of hydrogen and 2 atoms of oxygen

In molecular terms, 2 molecules of hydrogen combine with one molecule of oxygen to give two molecules of water.

The equation is now balanced!

Stoichiometry

Balancing a chemical equation is the first step to performing what chemists call Stoichiometry – using the balanced equation to determine amounts of reactants and products. It provides chemists, who have to deal with large quantities of actual substances, a way to understand the atomic world of atoms and molecules in terms of the real world of mass of substances.

To do this, chemists convert units of substances to Moles. Chemists have defined the mole of any substance in equivalent terms to the mass of atoms in 12 grams of Carbon-12. In other words, 1 mole of carbon weighs exactly 12 grams. Using comparative atomic masses between Carbon and other elements, the weight of other elements is determined. Hence, 1 mole of hydrogen weighs 1 gram, 1 mole of oxygen weighs 16 grams. And so on.

In stoichiometric terms the chemical equation to produce water, 2H₂(g)+ O₂(g)→ 2H₂O (l), is interpreted as follows – two moles of hydrogen react with one mole of oxygen to yield two moles of water. In other words, combining 2 grams of hydrogen with 16 grams of oxygen will yield 2 moles of water weighing 18 gm.

Check Point

1. One or more chemical substances, known as ______, chemically react to produce one or more ______.
2. “Energy can neither be created nor destroyed as a result of a chemical reaction”. This is the law of _________.
3. The number of _______ of each element in a chemical equation must remain the same on either side of that equation.
4. We make use of ______ to correctly balance chemical equations.
5. Iron (Fe) combines with oxygen (O₂) to form 2 units of iron oxide Fe₂O₃. Balance this equation: Fe (s) + O₂(g) →Fe₂O₃(s)

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