50 Examples Of Balanced Chemical Equations With Answers

50 Examples of Balanced Chemical Equations with Answers

Balancing chemical equations is a fundamental skill in chemistry. It ensures adherence to the law of conservation of mass, stating that matter cannot be created or destroyed in a chemical reaction. This means the number of atoms of each element must be the same on both sides (reactants and products) of the equation. This article provides 50 examples of balanced chemical equations, categorized for easier understanding, along with explanations to help you master this crucial concept.

Understanding Chemical Equations

Before diving into the examples, let's refresh our understanding of chemical equations. A chemical equation uses chemical formulas to represent a chemical reaction. The reactants (starting materials) are on the left side, and the products (resulting substances) are on the right side, separated by an arrow (→) indicating the direction of the reaction.

Example: H₂ + O₂ → H₂O (Unbalanced)

This equation represents the reaction between hydrogen (H₂) and oxygen (O₂) to produce water (H₂O). However, it's unbalanced because the number of oxygen atoms is not equal on both sides. Balancing requires adding coefficients (numbers placed before the formulas) to adjust the number of molecules involved.

Balanced Example: 2H₂ + O₂ → 2H₂O

Now, we have two hydrogen molecules reacting with one oxygen molecule to produce two water molecules. The number of hydrogen and oxygen atoms is equal on both sides (4 hydrogen and 2 oxygen atoms).

Categories of Balanced Chemical Equations

We'll categorize the examples to illustrate different reaction types and complexities:

I. Simple Combination Reactions (Synthesis)

Combination reactions involve two or more reactants combining to form a single product.

1. C + O₂ → CO₂ (Carbon reacting with oxygen to form carbon dioxide)
2. 2Na + Cl₂ → 2NaCl (Sodium reacting with chlorine to form sodium chloride)
3. Mg + S → MgS (Magnesium reacting with sulfur to form magnesium sulfide)
4. 2H₂ + O₂ → 2H₂O (Hydrogen reacting with oxygen to form water)
5. N₂ + 3H₂ → 2NH₃ (Nitrogen reacting with hydrogen to form ammonia)
6. CaO + H₂O → Ca(OH)₂ (Calcium oxide reacting with water to form calcium hydroxide)
7. SO₃ + H₂O → H₂SO₄ (Sulfur trioxide reacting with water to form sulfuric acid)
8. Fe + S → FeS (Iron reacting with sulfur to form iron sulfide)

II. Simple Decomposition Reactions

Decomposition reactions involve a single reactant breaking down into two or more simpler products.

9. 2HgO → 2Hg + O₂ (Mercury(II) oxide decomposing into mercury and oxygen)
10. 2KClO₃ → 2KCl + 3O₂ (Potassium chlorate decomposing into potassium chloride and oxygen)
11. CaCO₃ → CaO + CO₂ (Calcium carbonate decomposing into calcium oxide and carbon dioxide)
12. 2H₂O₂ → 2H₂O + O₂ (Hydrogen peroxide decomposing into water and oxygen)
13. 2NaCl → 2Na + Cl₂ (Sodium chloride decomposing into sodium and chlorine - requires high energy)

III. Single Displacement Reactions

Single displacement reactions involve one element replacing another in a compound.

14. Zn + 2HCl → ZnCl₂ + H₂ (Zinc replacing hydrogen in hydrochloric acid)
15. Fe + CuSO₄ → FeSO₄ + Cu (Iron replacing copper in copper sulfate)
16. 2Al + 3H₂SO₄ → Al₂(SO₄)₃ + 3H₂ (Aluminum replacing hydrogen in sulfuric acid)
17. Cu + 2AgNO₃ → Cu(NO₃)₂ + 2Ag (Copper replacing silver in silver nitrate)
18. Mg + 2H₂O → Mg(OH)₂ + H₂ (Magnesium reacting with water)

IV. Double Displacement Reactions (Metathesis)

Double displacement reactions involve the exchange of ions between two compounds.

19. AgNO₃ + NaCl → AgCl + NaNO₃ (Silver nitrate reacting with sodium chloride to form silver chloride and sodium nitrate)
20. Pb(NO₃)₂ + 2KI → PbI₂ + 2KNO₃ (Lead(II) nitrate reacting with potassium iodide to form lead(II) iodide and potassium nitrate)
21. BaCl₂ + Na₂SO₄ → BaSO₄ + 2NaCl (Barium chloride reacting with sodium sulfate to form barium sulfate and sodium chloride)
22. HCl + NaOH → NaCl + H₂O (Hydrochloric acid reacting with sodium hydroxide to form sodium chloride and water)
23. H₂SO₄ + 2KOH → K₂SO₄ + 2H₂O (Sulfuric acid reacting with potassium hydroxide to form potassium sulfate and water)
24. 2HNO₃ + Ca(OH)₂ → Ca(NO₃)₂ + 2H₂O (Nitric acid reacting with calcium hydroxide to form calcium nitrate and water)

V. Combustion Reactions

Combustion reactions involve a substance reacting with oxygen, often producing heat and light.

24. CH₄ + 2O₂ → CO₂ + 2H₂O (Methane burning in oxygen to form carbon dioxide and water)
25. C₂H₆ + 7/2O₂ → 2CO₂ + 3H₂O (Ethane burning in oxygen) Note: Fractional coefficients are sometimes used for simplicity
26. C₃H₈ + 5O₂ → 3CO₂ + 4H₂O (Propane burning in oxygen)
27. 2C₂H₂ + 5O₂ → 4CO₂ + 2H₂O (Acetylene burning in oxygen)
28. C₄H₁₀ + 13/2O₂ → 4CO₂ + 5H₂O (Butane burning in oxygen)

VI. Acid-Base Neutralization Reactions

These are a specific type of double displacement reaction involving an acid and a base. (Examples 21-23 above are also acid-base neutralization reactions)

29. 3HCl + Al(OH)₃ → AlCl₃ + 3H₂O (Hydrochloric acid neutralizing aluminum hydroxide)
30. H₂SO₄ + Mg(OH)₂ → MgSO₄ + 2H₂O (Sulfuric acid neutralizing magnesium hydroxide)

VII. More Complex Reactions

These examples showcase reactions involving multiple steps or more complex compounds.

31. 2Fe + 3Cl₂ → 2FeCl₃ (Iron reacting with chlorine to form iron(III) chloride)
32. 2K + 2H₂O → 2KOH + H₂ (Potassium reacting vigorously with water)
33. 4P + 5O₂ → 2P₂O₅ (Phosphorus burning in oxygen to form phosphorus pentoxide)
34. 2Al + Fe₂O₃ → Al₂O₃ + 2Fe (Thermite reaction: Aluminum reducing iron(III) oxide)
35. CH₄ + Cl₂ → CH₃Cl + HCl (Chlorination of methane) Note: This can lead to further chlorination

VIII. Reactions Involving Polyatomic Ions

36. 2NaOH + H₂SO₄ → Na₂SO₄ + 2H₂O (Sodium hydroxide reacting with sulfuric acid)
37. 3Ca(OH)₂ + 2H₃PO₄ → Ca₃(PO₄)₂ + 6H₂O (Calcium hydroxide reacting with phosphoric acid)
38. K₂CO₃ + 2HCl → 2KCl + H₂O + CO₂ (Potassium carbonate reacting with hydrochloric acid)
39. Na₂SO₃ + 2HCl → 2NaCl + H₂O + SO₂ (Sodium sulfite reacting with hydrochloric acid)

IX. Redox Reactions (Oxidation-Reduction)

Many of the above examples are also redox reactions, involving changes in oxidation states. Here are some more explicitly showing electron transfer:

40. 2Mg + O₂ → 2MgO (Magnesium oxidation)
41. Cu²⁺ + Zn → Cu + Zn²⁺ (Copper reduction by zinc)
42. Fe + CuSO₄ → FeSO₄ + Cu (Iron reduces copper(II) ions)

X. Precipitation Reactions

These form a solid precipitate as a product. (Examples 19 & 20 above are also precipitation reactions)

43. AgNO₃(aq) + KCl(aq) → AgCl(s) + KNO₃(aq) (Silver chloride precipitate forms)
44. BaCl₂(aq) + Na₂SO₄(aq) → BaSO₄(s) + 2NaCl(aq) (Barium sulfate precipitate forms)

XI. Gas Evolution Reactions

These produce a gaseous product. (Examples 24-28, 38, 39 above are gas evolution reactions)

45. Na₂CO₃(aq) + 2HCl(aq) → 2NaCl(aq) + H₂O(l) + CO₂(g) (Carbon dioxide gas evolves)
46. Zn(s) + 2HCl(aq) → ZnCl₂(aq) + H₂(g) (Hydrogen gas evolves)

XII. Organic Reactions (Simplified)

These reactions involve organic compounds (carbon-containing compounds).

47. C₂H₄ + H₂ → C₂H₆ (Ethene hydrogenation to ethane)
48. C₂H₅OH + 3O₂ → 2CO₂ + 3H₂O (Ethanol combustion)
49. CH₃COOH + NaOH → CH₃COONa + H₂O (Acetic acid neutralization)

XIII. Nuclear Reactions (Note: Balancing nuclear equations is different)

50. ²³⁵U + ¹n → ¹⁴¹Ba + ⁹²Kr + 3¹n (Nuclear fission of uranium) (Note: mass numbers and atomic numbers must balance separately)

This comprehensive list of 50 balanced chemical equations, categorized by reaction type, provides a solid foundation for understanding and practicing this critical skill in chemistry. Remember to always ensure the number of atoms of each element is the same on both sides of the equation. Practice regularly to improve your proficiency!