Chemical Equations: Reactants, Products, and Conservation of Mass

What a Chemical Equation Represents

A chemical equation is a compact way to describe a chemical reaction: which substances are used up (reactants), which substances are formed (products), and the relative amounts involved. It is not just a “recipe name”; it is a quantitative statement. When written correctly and balanced, a chemical equation communicates that atoms are rearranged, not created or destroyed, during the reaction.

A typical equation has this structure:

reactants  →  products

The arrow means “yields” or “produces.” Reactants are written on the left side of the arrow, products on the right. A plus sign separates multiple reactants or multiple products.

Example:

2 H2 + O2 → 2 H2O

This tells you that hydrogen gas and oxygen gas react to form water. The numbers in front (2, 1 implied, 2) are coefficients. Coefficients are essential because they encode the relative number of particles (molecules or formula units) participating in the reaction.

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Reactants vs. Products

Reactants are the starting materials. They are consumed as the reaction proceeds. Products are the substances formed. In many reactions, you can observe evidence of product formation: a gas bubbles out, a solid precipitate forms, a color changes, or heat/light is released or absorbed. These observations suggest a reaction occurred, but the equation tells you exactly what changed at the particle level: which atoms ended up in which new combinations.

What the Equation Does (and Does Not) Say

• It does say which substances are involved and the ratios in which they react and form.
• It does not automatically say how fast the reaction happens, how it is carried out, or whether it goes to completion. Those depend on conditions such as temperature, concentration, surface area, mixing, and catalysts.

Conservation of Mass and Why Equations Must Be Balanced

The conservation of mass means that in a closed system, the total mass of reactants equals the total mass of products. At the microscopic level, this is because atoms are conserved: chemical reactions rearrange atoms into new groupings, but they do not change one element into another. Therefore, a correct chemical equation must show the same number of atoms of each element on both sides.

Consider this unbalanced equation for forming water:

H2 + O2 → H2O

Count atoms:

• Left: H = 2, O = 2
• Right: H = 2, O = 1

Oxygen atoms do not match, so this violates conservation of mass. Balancing fixes the mismatch by adjusting coefficients (never by changing subscripts inside formulas). The balanced equation is:

2 H2 + O2 → 2 H2O

Now the atom counts match:

• Left: H = 4, O = 2
• Right: H = 4, O = 2

Why You Cannot Change Subscripts

Subscripts are part of the identity of a substance. Changing a subscript changes the substance itself. For example, H2O and H2O2 are different substances with different properties. Balancing is about changing how many units of each substance participate, not changing what the substances are.

Allowed:

2 H2O

Not allowed (because it changes the substance):

H2O2

Reading Coefficients as Ratios

Coefficients represent ratios of particles, and by extension ratios of moles. Even if you are not doing full mole calculations yet, it helps to interpret coefficients as “counts” of units.

In:

2 H2 + O2 → 2 H2O

• For every 2 molecules of H2, 1 molecule of O2 reacts.
• This produces 2 molecules of H2O.

The same ratio applies to moles:

• 2 moles H2 react with 1 mole O2 to produce 2 moles H2O.

It also applies to any proportional amount. If you had 200 molecules of H2, you would need 100 molecules of O2 to react completely, producing 200 molecules of H2O (assuming enough oxygen is available and the reaction goes to completion).

States of Matter and Reaction Conditions in Equations

Equations often include state symbols to show whether each substance is a solid, liquid, gas, or dissolved in water (aqueous). These symbols help you connect the equation to what you would see in a lab or everyday setting.

• (s) solid
• (l) liquid
• (g) gas
• (aq) aqueous (dissolved in water)

Example:

2 H2(g) + O2(g) → 2 H2O(l)

Sometimes conditions are written above or below the arrow, such as heat, light, or a catalyst. These are not reactants or products; they are conditions that affect the reaction.

2 KClO3(s) → 2 KCl(s) + 3 O2(g)

If a catalyst is used, it might be written above the arrow:

2 KClO3(s) → 2 KCl(s) + 3 O2(g)

(In a formatted textbook, the catalyst name would appear above the arrow; in plain text it may be noted nearby.) The key idea is that catalysts are not consumed; they help the reaction proceed.

Step-by-Step: How to Balance Chemical Equations

Balancing is a skill. The goal is to choose coefficients so that each element has the same number of atoms on both sides. A reliable approach is to balance one element at a time, using atom counts as your guide.

General Balancing Procedure

• Step 1: Write the correct formulas for all reactants and products. Do not guess; the formulas must represent the substances involved.
• Step 2: Count atoms of each element on both sides.
• Step 3: Add coefficients to make atom counts match. Start with elements that appear in only one reactant and one product when possible.
• Step 4: Recount after each change because changing one coefficient affects multiple elements.
• Step 5: Reduce coefficients to the smallest whole-number ratio.
• Step 6: Check your work by confirming every element matches.

Tip: If a polyatomic group (a set of atoms that stays together) appears unchanged on both sides, you can sometimes balance it as a unit. This is a shortcut that reduces errors, but always verify individual atom counts at the end.

Example 1: Formation of Ammonia

Unbalanced equation:

N2 + H2 → NH3

Step 1: Count atoms.

• Left: N = 2, H = 2
• Right: N = 1, H = 3

Step 2: Balance nitrogen first. Put a 2 in front of NH3 to make N = 2 on the right:

N2 + H2 → 2 NH3

Now count again:

• Left: N = 2, H = 2
• Right: N = 2, H = 6

Step 3: Balance hydrogen. Put a 3 in front of H2 to make H = 6 on the left:

N2 + 3 H2 → 2 NH3

Step 4: Check.

• Left: N = 2, H = 6
• Right: N = 2, H = 6

Balanced.

Example 2: Combustion of Propane (A Common Fuel)

Combustion reactions involve a fuel reacting with oxygen to produce carbon dioxide and water. For propane:

C3H8 + O2 → CO2 + H2O

Step 1: Balance carbon. There are 3 carbons on the left, so put 3 in front of CO2:

C3H8 + O2 → 3 CO2 + H2O

Step 2: Balance hydrogen. There are 8 hydrogens on the left, so put 4 in front of H2O (because 4 × 2 = 8):

C3H8 + O2 → 3 CO2 + 4 H2O

Step 3: Balance oxygen last. Now count oxygen atoms on the right:

• From 3 CO2: 3 × 2 = 6 O
• From 4 H2O: 4 × 1 = 4 O
• Total O on right = 10

On the left, oxygen comes as O2, so we need 10 oxygen atoms, which is 5 O2 molecules:

C3H8 + 5 O2 → 3 CO2 + 4 H2O

Step 4: Check. C: 3=3, H: 8=8, O: 10=10. Balanced.

Practical interpretation: the coefficients tell you that 1 unit of propane requires 5 units of oxygen and produces 3 units of carbon dioxide and 4 units of water (in particle or mole ratios).

Example 3: A Reaction That Produces a Precipitate

Many reactions in water form an insoluble solid called a precipitate. Consider:

NaCl(aq) + AgNO3(aq) → AgCl(s) + NaNO3(aq)

Step 1: Count atoms. Na: 1 both sides, Cl: 1 both sides, Ag: 1 both sides, N: 1 both sides, O: 3 both sides. It is already balanced with coefficients of 1.

Even when balancing is easy, the equation still communicates conservation of mass: the atoms present in dissolved ions are rearranged into a solid AgCl and a dissolved NaNO3.

Checking Conservation of Mass with a Simple Mass Example

Balanced equations support conservation of mass, but it can be helpful to see how this connects to measured masses. Suppose you react hydrogen and oxygen to form water:

2 H2 + O2 → 2 H2O

If you had a sealed container and you could measure masses precisely, the total mass before and after would match. For instance, if 4.0 g of hydrogen reacts completely with 32.0 g of oxygen, you would form 36.0 g of water (numbers chosen to illustrate the idea). The key point is not the specific values, but that mass is accounted for: reactant mass becomes product mass.

If the system is not closed, it may look like mass is not conserved. For example, if a gas escapes, the measured mass of the remaining container contents decreases. That is a measurement issue (mass leaving the system), not a violation of conservation of mass.

How to Identify Reactants and Products in Words and Convert to an Equation

Often you encounter reactions described in sentences. Your job is to identify the reactants and products and then write them in equation form. Focus on the structure of the sentence:

• “A reacts with B to form C and D.” → A + B → C + D
• “When A is heated, it decomposes into B and C.” → A → B + C

Example (word description): “Magnesium reacts with oxygen to form magnesium oxide.”

Mg + O2 → MgO

Now balance it:

2 Mg + O2 → 2 MgO

Step-by-step reasoning: oxygen is O2, but MgO contains one oxygen per unit, so you need 2 MgO to use both oxygen atoms, which then requires 2 Mg.

Common Patterns: Synthesis, Decomposition, Single Replacement, Double Replacement, Combustion

Recognizing reaction patterns helps you anticipate the form of products, which makes writing and balancing equations easier. The details of predicting products can be complex, but the patterns themselves are straightforward.

Synthesis (Combination)

Two or more substances combine to form one product.

A + B → AB

Example:

2 Mg + O2 → 2 MgO

Decomposition

One substance breaks apart into two or more products.

AB → A + B

Example:

2 H2O2 → 2 H2O + O2

Balancing note: oxygen appears in both products, so count carefully.

Single Replacement (Displacement)

One element replaces another in a compound.

A + BC → AC + B

Example:

Zn + 2 HCl → ZnCl2 + H2

Balancing note: H is diatomic as H2 in the product, so you need 2 HCl to supply 2 H atoms.

Double Replacement (Metathesis)

Two compounds exchange partners, often in water, sometimes forming a precipitate, gas, or water.

AB + CD → AD + CB

Example:

Na2SO4(aq) + BaCl2(aq) → BaSO4(s) + 2 NaCl(aq)

Step-by-step balancing: sulfate (SO4) stays together on both sides, so balance Na and Cl with coefficients. Left has 2 Na, so put 2 in front of NaCl.

Combustion

A substance reacts with oxygen, often producing CO2 and H2O when the fuel contains carbon and hydrogen.

fuel + O2 → CO2 + H2O

Example already balanced:

C3H8 + 5 O2 → 3 CO2 + 4 H2O

Step-by-Step: Balancing When You Get a Fraction

Sometimes balancing oxygen last leads to a fractional coefficient. This is acceptable as an intermediate step, but final coefficients should be whole numbers. You can clear fractions by multiplying every coefficient by the denominator.

Example: balance

Fe + O2 → Fe2O3

Step 1: Balance iron. Fe2O3 has 2 Fe, so put 2 in front of Fe:

2 Fe + O2 → Fe2O3

Step 2: Balance oxygen. Right has 3 O atoms; left has O2. To get 3 O atoms from O2, you would need 3/2 O2:

2 Fe + (3/2) O2 → Fe2O3

Step 3: Clear the fraction. Multiply every coefficient by 2:

4 Fe + 3 O2 → 2 Fe2O3

Step 4: Check. Fe: 4=4, O: 6=6. Balanced.

What It Means for a Reaction to “Go to Completion” and Why Excess Reactant Matters

Balanced equations give ratios, but real reactions depend on how much of each reactant you start with. Often one reactant runs out first; this reactant limits how much product can form. The other reactant is left over (in excess). Even without doing full calculations, you can reason using coefficients.

Using:

2 H2 + O2 → 2 H2O

If you start with 2 “units” of H2 and 1 “unit” of O2, they match the ratio exactly, so both can be fully consumed. If you start with 2 units of H2 and 2 units of O2, hydrogen will run out first, and oxygen will be left over. The equation helps you predict leftovers and product amounts qualitatively by comparing available amounts to the required ratio.

Common Mistakes and How to Avoid Them

• Changing subscripts instead of coefficients: Always keep formulas fixed; only adjust coefficients.
• Forgetting diatomic elements: Some elements commonly exist as two-atom molecules in their elemental form (for example, O2, N2, H2, Cl2). If you write them as single atoms when they are actually diatomic, balancing will not reflect the real substances.
• Balancing one element and not rechecking others: Every coefficient change affects atom counts; recount each time.
• Not reducing to simplest whole numbers: If all coefficients share a common factor (like 2), divide them to simplify.
• Ignoring states and context: Including (s), (l), (g), (aq) can prevent confusion, especially when a product is a gas or a precipitate.

Practice Workflow You Can Use on Any Equation

When you face a new reaction, use this repeatable workflow:

• Write the skeleton equation with correct formulas.
• Make a table of atom counts for each element on left and right.
• Balance elements that appear in the fewest places first.
• Leave oxygen and hydrogen for last in many cases, especially in combustion-type equations.
• If you get a fraction, clear it by multiplying all coefficients.
• Verify conservation of atoms for every element.

To make your checking systematic, you can write counts like this:

Element | Reactants | Products

Then fill in numbers and update them as you adjust coefficients. This simple habit prevents most balancing errors.