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 ↓


Introduction to Chemistry


Chemistry is a branch of science that studies the composition, structure, properties, and changes of matter. It attempts to understand the substances that make up our world and how they interact with each other.

What is the matter?

Matter is everything that has mass and occupies space. It includes everything you can see and touch, such as air, water, rocks, plants, animals, and even people. Matter can exist in different states, primarily known as solid, liquid, and gas. Each state has different properties:

  • Solid: It has a definite shape and volume.
  • Liquid: It has a definite volume, but it takes the shape of its container.
  • Gas: It has neither a definite shape nor a definite volume and it expands to fill its container.

Visualization of atoms in different states:

Solid Liquid Gas

Atoms and Molecules

At the core of chemistry is the concept of atoms, which are the basic building blocks of matter. An atom consists of a nucleus made up of protons and neutrons, surrounded by electrons.

Structure of the atom

An atom can be represented by this simple structure:

nucleus (protons + neutrons)
|
|
electrons

Molecules

Molecules are formed when two or more atoms join together. For example, a water molecule (H2O) is made up of two hydrogen atoms bonded to one oxygen atom.

Visual example of molecules

Here's a simplified representation of molecules:

H2O

Chemical reactions

When atoms and molecules interact, they undergo chemical reactions. These reactions result in changes in matter, often forming new substances with different properties. Chemical reactions are guided by the law of conservation of mass, which states that mass is neither created nor destroyed.

Chemical reactions are represented by chemical equations. For example:

2H2 + O2 → 2H2O

Visualization of reactants and products

The reactants and products in a chemical equation can be represented as follows:

H2 O2 H2O

Periodic table of elements

The periodic table is an important tool for chemists, serving as a comprehensive map of the known elements. Each element is placed on the table based on its atomic number, which indicates the number of protons in its nucleus. The table is arranged in rows called periods and columns called groups or families.

Structure of the periodic table

This table helps show trends in the properties of elements, such as electronegativities, ionization energies, and atomic radii. For example, elements in the same group often share chemical and physical properties.

Example element

Let's highlight some example elements:

Hydrogen (H) - Atomic Number: 1
Oxygen (O) - Atomic Number: 8
Carbon (C) - Atomic Number: 6
Sodium (Na) - Atomic Number: 11
Chlorine (Cl) - Atomic Number: 17

Basic Chemistry Concepts

Atomic number and mass number

The atomic number of an element is the number of protons in its nucleus. The mass number is the sum of the protons and neutrons in the nucleus. Consider carbon:

Carbon (C) Atomic number: 6
Mass number: 12 (6 protons + 6 neutrons)

Isotopes

Isotopes are atoms of the same element that have different numbers of neutrons. For example, hydrogen has three common isotopes:

  • Protium: 1 proton, 0 neutrons
  • Deuterium: 1 proton, 1 neutron
  • Tritium: 1 proton, 2 neutrons

Covalent and Ionic Bonding

Chemical bonds hold atoms together in molecules. There are two main types of chemical bonds:

  • Covalent bond: Atoms share electrons. For example, two hydrogen atoms share electrons to form H2 molecule.
  • Ionic bond: Atoms transfer electrons. An example of this is the bond formed between sodium and chloride in sodium chloride (table salt), NaCl.

Conclusion

Chemistry is a vast and fascinating field that affects every aspect of our daily lives. Understanding the basics of matter, atoms, molecules, and chemical reactions provides a valuable foundation for exploring more complex concepts in the science of chemistry. By studying the periodic table and fundamental chemical bonds, we appreciate the diversity of substances that make up the universe and how these substances interact to create what we see around us.


Undergraduate → 1.1


U
username
0%
completed in Undergraduate

← Prev (1)
General chemistry
Next (1.2) →
Atomic Structure

Comments 



Introduction to Chemistry

https://www.chemistryelite.com/ug/1.1?ceal=en