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 ↓


Atomic Structure


Understanding the structure of atoms is vital to the study of chemistry. Atoms are the basic building blocks of matter, and their structure determines the properties of the elements and compounds they form. In this lesson, we will explore the components of atoms, the historical developments that led to our current understanding, and the roles of subatomic particles.

Introduction to atoms

An atom is the smallest unit of an element that retains the chemical properties of that element. Atoms are made up of three primary types of particles: protons, neutrons, and electrons. Here is a simplified representation of an atom:

        ----------------------
        | Atom              |
        |                   |
        | Nucleus:          |
        | - Protons (p⁺)    |
        | - Neutrons (n⁰)   |
        |                   |
        | Electrons (e⁻)    |
        | orbiting nucleus  |
        ----------------------

Subatomic particles

Proton

Protons are positively charged particles located in the nucleus of an atom. The number of protons in the nucleus determines the atomic number and subsequently the identity of the element. For example:

  • Hydrogen has one proton: Atomic Number = 1
  • Helium has two protons: Atomic Number = 2
  • Carbon has six protons: Atomic Number = 6
  • Oxygen has eight protons: Atomic Number = 8

Neutron

Neutrons are neutral particles located in the nucleus along with the protons. They have no charge, and their primary role is to add mass to the atom and help stabilize the nucleus. The neutrons and protons together make up the mass number of the atom. The mass number is expressed as:

        Mass Number = Number of Protons + Number of Neutrons

Electrons

Electrons are negatively charged particles that orbit the nucleus. The number of electrons in an atom is equal to the number of protons in a neutral atom. Electrons affect the chemical properties of an atom, especially its ability to form bonds with other atoms. The distribution of electrons in an atom is arranged into energy levels or shells.

Historical perspective on atomic theory

The concept of the atom has evolved considerably over the centuries with contributions from many scientists. Some of the major developments are given below:

Democritus (c. 400 BC)

Democritus proposed that matter was composed of tiny, indivisible particles called "atomos," meaning indivisible. However, this early idea lacked experimental evidence.

John Dalton (1803)

John Dalton presented the first modern atomic theory, in which he proposed that atoms are indivisible particles that have different types depending on the element identity. His theories are as follows:

  • Elements are composed of tiny, indivisible particles called atoms.
  • All atoms of a given element are similar but different from the atoms of other elements.
  • Atoms can neither be created nor destroyed in chemical processes.

JJ Thomson (1897)

Thomson discovered the electron using cathode ray experiments. He proposed the "plum pudding model," which suggested atoms were spheres of positive charge with electrons embedded within them.

Ernest Rutherford (1911)

Rutherford's gold foil experiment demonstrated that atoms consist of a dense, positively charged nucleus surrounded by electrons, leading to the planetary model of the atom. This showed that most of the atom is empty space.

Niels Bohr (1913)

Bohr refined the atomic model by proposing that electrons travel in discrete orbits around the nucleus, and that electrons can jump between these orbits with quantized energy levels.

        Electron Transitions: n=3 ---> n=2 (Energy emitted as photon)

Quantum mechanical model

The current model of the atom is based on quantum mechanics, which provides a probabilistic view of the state of the electron within electron clouds or orbitals. This model is supported by wave functions that define the probability of the electron's position.

Nucleus electron cloud

The role of electrons in chemical bonding

The interaction of electrons facilitates chemical bonding. The electron configuration determines the way an atom can bond with other atoms:

Valence electrons

Valence electrons are the outermost electrons in an atom and are important in forming chemical bonds. Atoms achieve stable configurations (often like noble gases) through these electrons. For example:

  • Sodium (Na): 1s² 2s² 2p⁶ 3s¹ - 1 valence electron
  • Chlorine (Cl): 1s² 2s² 2p⁶ 3s² 3p⁵ - 7 valence electrons
  • Neon (Ne): 1s² 2s² 2p⁶ - 8 valence electrons (octet)

Atoms usually form covalent or ionic bonds to reach a stable electron configuration:

  • Ionic bonds: Usually formed between metals and nonmetals. Electrons are transferred, forming charged ions that attract each other.
  • Covalent bonds: Usually formed between nonmetals. Electrons are shared between atoms to achieve stability.

Isotopes and atomic mass

Isotopes are different forms of an element that have the same number of protons but different numbers of neutrons. This results in different mass numbers while retaining chemical properties:

  • Carbon-12 (C-12): 6 protons, 6 neutrons
  • Carbon-14 (C-14): 6 protons, 8 neutrons

The atomic mass of an element is the weighted average of the masses of its isotopes. The formula for average atomic mass is:

        Atomic Mass = Σ (Fraction of Isotope × Mass of Isotope)

Conclusion

The study of atomic structure is a cornerstone of chemistry, enabling a deeper understanding of the properties and behaviors of elements. Models of atomic structure have evolved over time, with the quantum mechanical model providing the most accurate description of atomic behavior. Understanding subatomic particles and their interactions helps to understand chemical bonding, isotope behavior, and atomic mass, which are important in all areas of chemistry.


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Atomic Structure

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