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 → Stoichiometry ↓


Mole concept


The mole concept is one of the cornerstones of the study of chemistry. It is a fundamental concept that allows chemists to calculate and measure substances on the atomic and molecular scale, which is otherwise difficult due to the small size of atoms and molecules. By learning and applying the mole concept, we can solve a wide range of chemical problems and understand the stoichiometry of chemical reactions. This lesson will help you understand the mole concept and its applications in chemistry in depth, including definitions, calculations, and examples.

What is a mole?

The concept of "mole" helps chemists keep track of units such as atoms, molecules, ions, and other small particles when working with large quantities. A mole is similar to a unit such as a dozen, which simply means 12 items. In chemistry, a mole is defined as Avogadro's number of particles, which is about 6.022 x 10 23 particles.

No. of particles in a mole = 6.022 x 10^23

This number is named after Amedeo Avogadro, an Italian scientist who contributed to the molecular theory in chemistry.

Understanding Avogadro's number

1 mole = 6.022 x 10 23 particles

Avogadro's number, 6.022 x 10 23, is important because it establishes a direct connection between macroscopic and microscopic scales. When you need to adjust from atomic-scale to practical laboratory measurements, the mole and Avogadro's number make this possible.

Molar mass

Molar mass is the mass of one mole of a given substance (chemical element or chemical compound) and is expressed in grams per mole (g/mol). It provides a bridge between the mass of a substance and the number of moles of that substance. To find the molar mass, you simply need to add up the atomic masses of all the atoms in a molecule or formula unit, as given in the periodic table.

For example, let's calculate the molar mass of water (H 2 O):

  1. The molar mass of hydrogen (H) is about 1 g/mol.
  2. The molar mass of oxygen (O) is about 16 g/mol.
  3. Total mass: 2(1 g/mol) + 16 g/mol = 18 g/mol

The molar mass of water is about 18 g/mol.

Converting between moles and particles

Converting between moles and the number of particles (atoms, molecules, etc.) becomes simple once you understand the use of Avogadro's number. The formula for this conversion is:

No. of particles = No. of moles x 6.022 x 10^23

For example, if you have 3 moles of oxygen molecules, the number of molecules will be:

No. of O 2 molecules = 3 moles x 6.022 x 10^23 molecules/mole = 1.806 x 10^24 molecules

Converting between moles and mass

Conversion between mass and number of moles of a substance involves using molar mass. The conversion formulas are:

No. of moles = Mass of substance (g) / Molar Mass (g/mol)
Mass of substance (g) = No. of moles x Molar Mass (g/mol)

Suppose you have 36 grams of water. To find the number of moles:

No. of moles = 36 g / 18 g/mol = 2 moles

Example problems

Example 1: Atoms in one mole of a sample

Calculate how many atoms are in 2 moles of aluminum (Al).

No. of Atoms = No. of Moles x Avogadro's Number = 2 moles x 6.022 x 10^23 atoms/mole = 1.2044 x 10^24 atoms

Thus, 2 moles contain 1.2044 x 10 24 atoms of aluminum.

Example 2: Mole to mass conversion

If you have 0.5 moles of carbon dioxide (CO 2), calculate its mass.

First, calculate the molar mass of CO2:

  • Carbon: 12 g/mol
  • Oxygen: 16 g/mol each
  • Total: 12 g/mol + 2(16 g/mol) = 44 g/mol

Mass = number of moles x molar mass:

Mass = 0.5 moles x 44 g/mol = 22 grams

Therefore, 0.5 moles of CO2 is 22 grams.

Use of mole concept in stoichiometry

Stoichiometry is the field of chemistry that involves calculating the amounts of reactants and products in chemical reactions. The mole concept is fundamental to stoichiometry because it allows masses to be converted into quantities of atoms and molecules where direct calculations are impossible. With balanced chemical equations, stoichiometric calculations can determine the amounts of substances consumed and produced.

Balancing chemical equations

Take the simple combustion reaction of methane (CH4):

CH 4 + 2O 2 → CO 2 + 2H 2 O

This balanced equation shows that one mole of methane reacts with two moles of oxygen to form one mole of carbon dioxide and two moles of water.

Stoichiometric calculations

Suppose we want to calculate how much water will be formed from 5 moles of methane, using the balanced equation.

Based on equation: 1 mole CH 4 produces 2 moles H 2 O Therefore: 5 moles CH 4 produces 5 x 2 = 10 moles H 2 O

Thus, 5 moles of methane will produce 10 moles of water.

Limitations of the mole concept

Despite being effective, the mole concept can have some limitations. The assumption of ideal conditions is rarely true outside of theoretical or controlled laboratory settings. In real-world chemical reactions, factors such as reaction rate, temperature, and pressure can affect results beyond theoretical calculations.

The role of the mole concept in chemistry

The concept of the mole is fundamental to chemistry because it connects the microscopic world of atoms and molecules to the macroscopic world that can be observed in the laboratory. By converting atoms and molecules into moles, chemists can precisely understand reactions and the behavior of compounds. This concept enables the production of medicines in the laboratory, the development of new materials, and many other essential applications in science and technology.

Conclusion

Understanding the mole concept provides important capabilities for studying and applying chemistry outside of symbolic equations and theoretical confines. By learning to convert and calculate between volume, mass, and particles, chemists can predict and design reactions needed in scientific research and industry applications. The mole concept will continue to be an important standard in understanding new modern chemical insights and developing practical solutions in the future.


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Mole concept

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