Atomic Theory Flashcards

Atomic Theory: From Greeks to Present

The Atom - A Tiny Bit From History

• The study of the atom traces back to early philosophical ideas about matter.

The Early Greeks

• Aristotle (4th century BC): Proposed that matter had four properties:

  • Moist

  • Dry

  • Hot

  • Cold

• These properties were contained in various proportions by four major elements:

  • Fire

  • Air

  • Earth

  • Water

• Early Greek Periodic Table (Hypothetical):

  • Air: Moist, Water

  • Fire: Hot, Cold

  • Earth: Dry

Medieval Chemists

• Developed the idea of "corpuscles" subject to attractive and repulsive forces.

• Ideas from Greeks and medieval chemists were mostly philosophical, not based on the scientific method.

Dalton’s Atomic Theory (1808)

• John Dalton: Based his model on solid experimental discoveries.

• Key hypotheses:

  • Elements are made up of extremely small particles called atoms.

  • All atoms of a given element are identical.

  • Each compound is unique and consists of particular atoms arranged in a particular way.

  • Chemical reactions involve reshuffling of atoms to form new compounds from the old atoms.

• Dalton's Laws:

  • Law of Definite Proportions: A chemical compound always contains exactly the same proportion of elements by mass.

  • Law of Multiple Proportions: When elements combine, they do so in the ratio of small whole numbers.

  • Law of Conservation of Mass: The mass of a closed system remains constant, regardless of the processes acting inside the system.

The Thompson Atom

• J.J. Thompson (Mid-19th Century): Discovered and demonstrated that atoms have negatively charged "electrons" and positively charged "protons."

• "Plum Pudding" Model:

  • Atoms consist of negative electrons scattered within a positive "pudding."

The Rutherford Atom Model

• Rutherford's Proposal:

  • The atom consists of a tiny, positively charged nucleus surrounded by a "cloud" of negatively charged electrons.

  • The number of protons is equal to the number of electrons, making the atom electrically neutral.

• Gold Foil Experiment:

  • Rutherford's Gold Foil Experiment proved the existence of a small massive center to atoms, which would later be known as the nucleus of an atom.

  • The majority of alpha particles passed straight through the gold foil, with only a small number reflecting violently.

Chadwick

• James Chadwick (1932): Discovered the neutron, which had been predicted by Rutherford.

Bohr's Model

• Niels Bohr: Expanded on Rutherford’s model.

• Key Points:

  • Electrons are restricted to specific paths called "orbitals" at fixed distances from the nucleus.

  • Electrons can only emit or absorb energy when they move from one orbital to another.

  • Only certain quantities of energy are emitted or absorbed (quantized energy levels).

Bohr Diagrams

• Simplified pictures showing the arrangement of electrons in atoms and ions.

• Electrons are located in circular paths (orbitals) around the nucleus.

• Orbitals can contain up to 2, 8, 8, 18, 18 electrons.

• Valence electrons are those in the outermost orbital (bonding electrons).

• Steps to write Bohr Diagrams:

  1. Find the number of protons, neutrons, and electrons.

  2. Write "P=" and "N=" for the number of protons and neutrons in the nucleus.

  3. Draw circles around the "nucleus" and fill them with electrons in the pattern 2, 8, 8, 18, 18 until all electrons are accounted for.

Key Relationships

• # of protons = Atomic number

• # of neutrons = Atomic Mass – Atomic number

• # of electrons = # of protons for a neutral atom

• # of electrons = # of protons – (charge) for an ion

Bohr Diagram Examples

• Fluorine atom example (9p, 10n)

• Fluoride ion (F-) example (9p, 10n)

Atomic Number and Atomic Mass

• Elements are differentiated by the number of protons in the nucleus.

• Atomic Number: The number of protons in the nucleus.

• In a neutral atom: Atomic Number = Number of Electrons = Number of Protons.

• Information on the Periodic Table:

  • Atomic Number (whole number)

  • Atomic Symbol

  • Atomic Mass (decimal)

• Neutrons and protons have a mass of approximately 1.0 amu, so the total atomic mass is found by their combined totals.

Electrons and Ions

• The mass of an electron is approximately 0 amu.

• Ion: A particle formed when electrons are added or subtracted from an atom.

• Calculating Electrons in Ions: #e- = atomic # – (charge)

  • Example a) N^{3-}: #e = 7 - (-3) = 10 e

  • Example b) Ca^{2+}: #e = 20 - (+2) = 18 e

Atomic Particles Examples

• Find the number of protons, neutrons, and electrons:

  • a) ^{27.0}_{13}Al: p = 13, n = 14, e = 13

  • b) ^{74.9}_{33}As: p = 33, n = 42, e = 33

Isotopes

• Species with the same atomic number but different atomic masses (same # of protons, different # of neutrons).

• Examples:

  • ^1_1H = H = ordinary Hydrogen (protium)

  • ^2_1H = D = Deuterium (heavy hydrogen)

  • ^3_1H = T = Tritium (radioactive hydrogen)

Molar Masses

• Molar masses on the periodic table are average masses of a sample containing a mixture of isotopes.

• Example (Carbon):

  • C-12 (97.00%), C-13 (1.00%), C-14 (2.00%)

  • C-12: 0.97 x 12 g/mol = 11.64 g/mol

  • C-13: 0.01 x 13 g/mol = 0.13 g/mol

  • C-14: 0.02 x 14 g/mol = 0.28 g/mol

  • Total: 12.05 g/mol

Molar Mass Examples Continued

• Chlorine: 75.77% Cl-35 and 24.23% Cl-37

  • Cl-35: 0.7577 x 34.966653 g/mol = 26.49423298 g/mol

  • Cl-37: 0.2423 x 36.965903 g/mol = 8.95683829 g/mol

  • Average molar mass: 35.45 g/mol