Undergraduate
  1. General chemistry  1.1. Introduction to Chemistry  1.2. Atomic Structure  1.2.1. Subatomic particles  1.2.2. Atomic Model  1.2.3. Quantum numbers  1.2.4. Electron Configuration  1.2.5. Periodic trends  1.2.6. Isotopes and their applications  1.3. Chemical bond  1.3.1. Ionic bond  1.3.2. Covalent bonds  1.3.3. Metal Bond  1.3.4. Hybridization  1.3.5. Molecular geometry and VSEPR theory  1.3.6. Bond polarity and dipole moment  1.4. Stoichiometry  1.4.1. Mole concept  1.4.2. Empirical and molecular formula  1.4.3. Limiting reactant and excess reactant  1.4.4. Percent Yield and Purity  1.4.5. Dilution calculation  1.5. States of matter  1.5.1. Gas Laws  1.5.2. Matter and Intermolecular Forces  1.5.3. Solid and Crystal Structures  1.5.4. Phase diagram  1.5.5. Plasma and supercritical fluid  1.6. Chemical reactions  1.6.1. Types of Chemical Reactions  1.6.2. Balancing Chemical Equations  1.6.3. Thermochemical reactions  1.6.4. Reaction energy profiles  1.7. Solutions and Mixtures  1.7.1. Concentration units  1.7.2. Syndrome properties  1.7.3. Solubility and Solubility Rules  1.7.4. Henry's Law and Raoult's Law  1.7.5. Partition coefficient  1.8. Chemical equilibrium  1.8.1. Law of mass action  1.8.2. Le Chatelier's principle  1.8.3. Reaction quotient  1.8.4. To use this online calculator for Equilibrium Constant, enter Equilibrium Constant (Eq.) and hit the calculate button. Here is how the Equilibrium Constant calculation can be explained with given input values -> 0.05 = 0.05/0.  1.8.5. Acid-base balance  1.9. Acids and bases  1.9.1. Arrhenius, Brønsted–Lowry, and Lewis definitions  1.9.2. pH and pOH  1.9.3. Acid-base titration  1.9.4. Buffer Solution  1.9.5. Acid-base indicator  1.9.6. Polyprotic Acids and Bases  1.10. Electrochemistry  1.10.1. redox reactions  1.10.2. Galvanic and Electrolytic Cells  1.10.3. Nernst equation  1.10.4. Electrolysis  1.10.5. Faraday's laws of electrolysis  1.11. Thermodynamics  1.11.1. Laws of Thermodynamics  1.11.2. Enthalpy, entropy and Gibbs free energy  1.11.3. Spontaneity of reactions  1.11.4. Thermodynamic cycles (Hess's law, Carnot cycle)  1.12. Kinetics  1.12.1. Rate law and reaction order  1.12.2. Activation Energy and Catalyst  1.12.3. Collision theory  1.12.4. Reaction mechanism  1.12.5. Transition state theory
  2. Organic chemistry  2.1. Structure and relationships  2.1.1. Hybridisation in Carbon Compounds  2.1.2. Resonance and aromatization  2.1.3. Molecular orbitals  2.1.4. Inductive and mesomeric effects  2.2. Hydrocarbons  2.2.1. Hydrocarbons  2.2.2. Alkene  2.2.3. Alkynes  2.2.4. Aromatic compounds  2.2.5. Cycloalkanes and Conformation  2.3. Functional Group  2.3.1. Alcohols and phenols  2.3.2. Ethers and epoxides  2.3.3. Aldehyde and ketone  2.3.4. Carboxylic acids and derivatives  2.3.5. Amin  2.3.6. Esters and amides  2.3.7. Thiols and sulfides  2.4. Stereoscopic  2.4.1. Isomerism in Stereochemistry  2.4.2. Chirality and optical activity  2.4.3. Structural analysis  2.4.4. Diastereomers and Enantiomers  2.5.1. Substitution reactions  2.5.2. Elimination reactions  2.5.3. Addition reactions  2.5.4. Radical reactions  2.5.5. Rearrangement reactions  2.5.6. Pericyclic Reactions  2.6. Spectroscopy and structural analysis  2.6.1. Infrared Spectroscopy  2.6.2. Nuclear magnetic resonance spectroscopy  2.6.3. Mass Spectrometry  2.6.4. UV-visible spectroscopy  2.6.5. X-ray crystallography  2.7. Polymer chemistry  2.7.1. Addition and Condensation Polymers  2.7.2. Polymerization mechanism  2.7.3. Biodegradable Polymer  2.7.4. Conducting and smart polymers
  3. Inorganic chemistry  3.1. Coordination chemistry  3.1.1. Ligands and coordination compounds  3.1.2. Crystal field theory  3.1.3. Molecular Orbital Theory in Coordination Compounds  3.1.4. Jahn–Taylor distortion  3.2. Main Group Chemistry  3.2.1. Alkali and alkaline earth metals  3.2.2. Halogens and Noble Gases  3.2.3. Transition metals and their complexes  3.2.4. Lanthanides and Actinides  3.3. Solid state chemistry  3.3.1. Crystal lattices and unit cells  3.3.2. Defects in solids  3.3.3. Electrical and magnetic properties  3.3.4. Superconductivity  3.4. Bio-inorganic chemistry  3.4.1. Metal ions in biology  3.4.2. Enzymatic reactions with metals  3.4.3. Metal poisoning and detoxification  3.4.4. Metalloproteins and Metalloenzymes
  4. Physical chemistry  4.1. Quantum Chemistry  4.1.1. Wave–particle duality  4.1.2. Schrödinger Equation  4.1.3. Quantum Numbers and Orbitals  4.1.4. Atomic and molecular spectroscopy  4.2. Statistical mechanics  4.2.1. Maxwell–Boltzmann distribution  4.2.2. Partitioning functions  4.2.3. Bose–Einstein and Fermi–Dirac statistics  4.3. Chemical thermodynamics  4.3.1. Gibbs Free Energy  4.3.2. Phase transition  4.3.3. Thermodynamic efficiency  4.4. Surface chemistry  4.4.1. Absorption and Catalysis  4.4.2. Colloids and Emulsions
  5. Analytical Chemistry  5.1. Classical methods  5.1.1. Gravimetric Analysis  5.1.2. titrimeter  5.2. Instrumental methods  5.2.1. Chromatography  5.2.2. Spectroscopy  5.2.3. Electroanalytical methods  5.2.4. X-ray diffraction
  6. Industrial Chemistry  6.1. Petrochemistry  6.2. Chemical engineering principles  6.3. Green Chemistry  6.4. Pharmaceuticals and Drug Synthesis
  7. Biochemistry  7.1. Carbohydrates  7.2. Proteins and enzymes  7.3. Lipids and membranes  7.4. Nucleic acids  7.5. Metabolism and Bioenergetics  7.6. Enzyme kinetics and mechanism

Undergraduate → General chemistry → Chemical reactions ↓


Balancing Chemical Equations


In the world of chemistry, balancing chemical equations is a basic skill. It is necessary to ensure that a chemical reaction follows the law of conservation of mass, which states that matter cannot be created or destroyed in an isolated system. This means that the number of atoms for each element must be the same on both sides of a chemical equation.

Understanding chemical equations

A chemical equation is a written representation of a chemical reaction. It shows reactants (starting substances) changing into products (substances formed). For example, consider hydrogen gas reacting with oxygen gas to form water:

H 2 + O 2 → H 2 O

In this equation, H 2 and O 2 are the reactants, and H 2 O is the product.

Law of conservation of mass

As mentioned, the law of conservation of mass dictates that in a chemical reaction, the mass of the products must be equal to the mass of the reactants. Therefore, the number of atoms of each element must be the same on both sides of the equation. When this condition is met, the equation is considered "balanced."

Steps to balance a chemical equation

Balancing chemical equations involves a step-by-step approach:

  1. Identify the reactants and products. Write the unbalanced equation using the correct chemical formulas.
  2. Count the number of atoms of each element: For each element, count the number of atoms on both the reactant and product sides.
  3. Add coefficients to balance atoms. Add whole number coefficients in front of chemical formulas so that there are the same number of atoms of each element on both sides.
  4. Repeat if necessary. Check and rebalance if necessary until all elements are balanced.

Example: Balancing a simple chemical equation

Let's balance the equation for the reaction of hydrogen gas with oxygen gas to form water:

Unbalanced: H 2 + O 2 → H 2 O

Start by counting the atoms:

  • Hydrogen (H): 2 on the left, 2 on the right (H 2 O has 1 H 2 )
  • Oxygen (O): 2 on the left, 1 on the right (H 2 O has 1 O)

The oxygen atoms are unbalanced. To balance them we need two water molecules:

H 2 + O 2 → 2 H 2 O

Now, count the atoms:

  • Hydrogen (H): 2 on the left, 4 on the right
  • Oxygen (O): 2 on the left, 2 on the right

The hydrogens are now unbalanced. To correct this, add a coefficient of 2 next to H 2 :

2 H 2 + O 2 → 2 H 2 O

Now the equation is balanced, there are 4 hydrogen atoms and 2 oxygen atoms on each side.

Visual example: Balancing equations using symbols

H2 O2 H2O

In the SVG above, a simplified schematic view shows the balance: two blue rectangles represent H 2 and a red rectangle represents O 2, while the gray rectangle corresponds to H 2 O Adjustments to the coefficients are necessary to balance this illustration, similar to the equation. For example, adding a second gray rectangle would indicate the need for two water molecules.

Advanced example: Balancing a more complex chemical equation

Let's balance a slightly more complicated reaction: the combustion of propane (C 3 H 8 ) in oxygen, forming carbon dioxide and water:

C 3 H 8 + O 2 → CO 2 + H 2 O

Count the atoms of each element:

  • Carbon (C): 3 on the left, 1 on the right
  • Hydrogen (H): 8 on the left, 2 on the right
  • Oxygen (O): 2 on the left, 3 on the right

Balance the carbon by placing a coefficient of 3 in front of CO 2 :

C 3 H 8 + O 2 → 3 CO 2 + H 2 O

Now balance the hydrogen by placing a coefficient of 4 in front of H 2 O :

C 3 H 8 + O 2 → 3 CO 2 + 4 H 2 O

Count the oxygen atoms again. There are (3 × 2) + 4 = 10 oxygen atoms on the right side. Put a coefficient of 5 in front of O 2 on the left side for balance:

C 3 H 8 + 5 O 2 → 3 CO 2 + 4 H 2 O

Final verification shows that the equation is balanced, with 3 carbons, 8 hydrogens, and 10 oxygens on both sides.

Checking your work

After balancing a chemical equation, it is important to verify accuracy by recalculating the atoms of each element to ensure equality on both sides. Mistakes can often occur with complex equations if coefficients are incorrect or omitted.

Common mistakes in balancing equations

  • Fractions in coefficients: It is important to avoid using fractions. Use whole number coefficients and adjust them if necessary by multiplying them by the least common multiple.
  • Ignoring polyatomic ions: If polyatomic ions appear unchanged on both sides of the equation, treat them as a single unit.
  • Ignoring the Law of Conservation: Always remember that balancing equations involves making sure that each type of atom is conserved during the transformation.

Tips for effective balancing

  • Balance one element at a time: Address one element fully before moving on to the next. A systematic approach helps with complex equations.
  • Leave O and H for last: Often, it's best to balance the oxygen and hydrogen atoms after all other elements.
  • Adjust polyatomic ions together: If a polyatomic ion remains unchanged during the reaction, treat it as a single entity for counting and balancing.

Practice problems

Try balancing the following equations to enhance your understanding:

  1. Fe + O 2 → Fe 2 O 3
  2. CaCl 2 + AgNO 3 → Ca(NO 3 ) 2 + AgCl
  3. (NH 4 ) 2 SO 4 + NaOH → Na 2 SO 4 + NH 3 + H 2 O
  4. C 4 H 10 + O 2 → CO 2 + H 2 O

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Balancing Chemical Equations

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