1. States of Matter. The Atom and the Periodic Table
1.4. Atomic Theory. The Atomic Model
The atomic theory
Democritus first suggested that matter may be made of indivisible particles = atoms. John Dalton provided the first scientific atomic theory (based on experiments) supporting the idea of "atoms" (1805).
Dalton's atomic theory stated that:
- Matter is made of solid "marble-like" particles (atoms).
- The mass of the atom decides what the properties of the atoms are.
Thanks to Dalton's new theory, elements can be said to be made of atoms. Jöns Jacob Berzelius compiles a table with the relative masses of the then-known atoms (1818), which are later used by Dmitri Mendeleev to organize the periodic table (1869).
Studies uranium salts (1896).
- Discovers that a photographic plate is exposed by the uranium.
- Not caused by x-rays!
Characterizes the radiation from e.g. radium (1896).
- Coins the term ”radioactivity”.
- The atom is not indivisible!
Nobel prize in physics 1903 (together with her husband Pierre, and Henri Becquerel).
J. J. Thomson
Uses a cathode ray tube to discover and characterize the electron (1898).
The "plum pudding" model
The electrons are spread throughout the atom, like plums in a plum pudding.
The gold foil experiment
Alpha particles beamed at a thin gold foil.
- Thomsons "plum pudding" model predicted: Particles should pass right through.
- Reality: About 1/8000 particles bounced of a very dense, solid nucleus.
Rutherford's atomic model
- A very small, dense, and positively charged nucleus.
- Electrons circle around, like ”planets”.
Problem with Rutherford's model: The electrons should immediately fall back into the nucleus.
- Bohr studied the emission spectrum from hydrogen.
- Only certain colors (energies) are emitted ⇒ The electrons may only reside in certain energy levels (electron "shells").
Beyond Bohr's model
- In reality, the electron shells are not that simple.
- The electrons are not negatively charged ”marbles”, but rather ”smeared” in time and space.
- Electron clouds rather than electron shells.
Protons and neutrons
Protons discovered by Rutherford (1917).
- Problem: How can the protons stay in the nucleus?
- There must be some kind of ”glue”!
Neutrons discovered by Rutherford's student, James Chadwick (1932).
Summary: The building blocks of the atom
- Charge: –1 (e–)
- In different energy levels (shells) around the nucleus
- Charge: +1 (p+)
- In the nucleus
- Charge: 0 (n)
- In the nucleus
- 1. States of Matter. The Atom and the Periodic Table
- 1.1. Matter. States of Matter
- 1.2. Elements and Chemical Compounds. Pure Substances and Mixtures
- 1.3. The Birth of Chemistry
- 1.4. Atomic Theory. The Atomic Model
- 1.5. Atomic Number, Mass Number, and Atomic Mass
- 1.6. Electron Configurations
- 1.7. Beyond Bohr's Atomic Model
- 1.8. Redox Reactions
- 1.9. The Structure of the Periodic Table
- 1.10. The Noble Gases
- 1.11. The Alkali Metals and the Halogens
- 1.12. The Alkaline Earth Metals and the Oxygen Group
- 1.13. A Few of the Elements in Group 13, 14, and 15
- 2. Chemical Calculations
- 2.1. Physical Quantity, Magnitude, and Units
- 2.2. Atomic Mass, Molecular Mass, and Unit Mass
- 2.3. Amount of Substance, Molar Mass, and Mass
- 2.4. Stoichiometry. Conservation of mass
- 2.5. Water of Crystallization
- 2.6. Calculating the Formula of a Chemical Compound
- 2.7. From Empirical to Molecular Formulas
- 2.8. Equivalent Amounts of Substance and Masses
- 2.9. Gases and Pressure
- 2.10. Concentrations
- 2.11. Dilutions
- 2.12. Yield
- 2.13. Limiting Reactants
- 3. Chemical Bonding
- 3.1. How Ionic Compounds are Formed
- 3.2. Precipitations
- 3.3. Names and Formulas of Ionic Compounds
- 3.4. Ionic Bonds
- 3.5. Properties of Ionic Compounds
- 3.6. Metal Bonding
- 3.7. Covalent Bonding
- 3.8. Polar Covalent Bonding
- 3.9. Dipoles. Polar and non-polar Molecules
- 3.10. The VSEPR Theory
- 3.11. Hydrogen Bonding. The Peculiar Water
- 3.12. Equals Solves Equal
- 3.13. Solubility of Gases in Water
- 3.14. Solubility of Salts in Water
- 4. Thermochemistry
- 5. Chemical Equilibrium
- 5.1. Reaction Rates
- 5.2. The Law of Mass Action
- 5.3. Calculations on Chemical Equilibrium
- 5.4. Heterogenous Equilibria. Solubility Product
- 5.5. Is the System at Equilibrium? The Reaction Quotient Q
- 5.6. Changing the Concentrations in a System in Equilibrium.
- 5.7. Diluting or Compressing Systems in Equilibrium, or Changing the Temperature
- 6. Acids and bases
- 7. Oxidation and Reduction
- 8. Electrochemistry
- 9. Organic Chemistry
- 9.1. Alkanes
- 9.2. Chain Isomers. Nomenclature
- 9.3. Haloalkanes
- 9.4. Nucleophilic Substitution
- 9.5. Alkenes
- 9.6. Electrophilic Addition. Markovnikov’s Rule
- 9.7. Elimination
- 9.8. Alkynes
- 9.9. Arenes and Aromatic Compounds
- 9.10. Alcohols
- 9.11. Oxidation of Alcohols
- 9.12. Aldehydes and Ketones
- 9.13. Thiols and Disulfides
- 9.14. Ethers
- 9.15. Amines
- 9.16. Nitro Compounds and Organic Nitrates
- 9.17. Carboxylic Acids
- 9.18. More on Carboxylic Acids
- 9.19. Stereoisomerism
- 9.20. Esters
- 9.21. Lipids
- 9.22. Mono-, Oligo-, and Polysaccharides
- 9.23. Amino Acids
- 9.24. Nucleotides
- 10. Biochemistry
- 10.1. Proteins
- 10.2. Enzymes
- 10.3. Catabolic Processes
- 10.4. Carrier Molecules
- 10.5. Glycolysis
- 10.6. Beta-oxidation
- 10.7. The Citric Acid Cycle
- 10.8. The Metabolism of Amino Acids
- 10.9. The Electron Transport Chain
- 10.10. Anabolic Processes
- 10.11. Gluconeogenesis and Fatty Acid Synthesis
- 10.12. DNA: Structure and Function
- 11. Analytical chemistry