History of the Atomic Model

Early Greek Theories Aristotle (350 B.C.) 4 Elements Democritus (400 B.C) Atoms and a void (empty space) Atoms are indivisible Based on philosophical arguments, NOT experimental evidence

Dalton’s Billiard Ball Model (1805) All matter is made of tiny, indivisible particles called atoms. Each element is made up of atoms that are unique and unlike atoms of any other element. Atoms of an element are identical. Matter is composed of indestructible, indivisible atoms

Thomson’s Raisin Bun Model (1897) Materials, when rubbed, can develop a charge difference. This electricity was called “cathode rays” These rays have a small mass and are negatively charged. Thomson noted that these negative subatomic particles (electrons) were a fundamental part of all atoms.

Rutherford’s Nuclear Model (1911) Rutherford shot alpha () particles at gold foil. Zinc sulfide screen Thin gold foil Lead block Radioactive substance path of invisible -particles Most particles passed through. So, atoms are mostly empty space. Some positive -particles deflected or bounced back! Thus, a “nucleus” is positive (protons) & holds most of an atom’s mass.

Table 1, p. 26

Practice P. 26 #1-5

Atomic numbers, Mass numbers Elements are often symbolized with their mass number (A) and atomic number (Z) E.g. Oxygen: O 16 8 Z = # of protons = # of electrons A - Z = # of neutrons Calculate # of e–, n0, p+ for Ca, Ar, and Br

Atomic Mass p+ e- no Ca 20 40 20 20 20 Ar 18 40 18 18 22 Br 35 80 35 35 45

Bohr - Rutherford diagrams Putting all this together, we get B-R diagrams To draw them you must know the # of protons, neutrons, and electrons (2,8,8,18 filling order) Draw protons (p+), (n0) in circle (i.e. “nucleus”) Draw electrons around in shells 3 p+ 4 n0 2e– 1e– Li shorthand 2 p+ 2 n0 He Li Draw Be, B, Al and shorthand diagrams for O, Na

Be B Al O Na 8 p+ 11 p+ 8 n° 12 n° 4 p+ 5 n° 5 p+ 6 n° 13 p+ 14 n° 2e– 8e– 1e– Na 8 p+ 8 n° 2e– 6e– O

Isotopes and Radioisotopes Isotopes: Atoms of the same element that have different numbers of neutrons Due to isotopes, mass #s are not round #s. E.g. Li (6.9) is made up of both 6Li and 7Li. Often, at least one isotope is unstable.It breaks down, releasing radioactivity.These types of isotopes are called radioisotopes Q- Sometimes an isotope is written without its atomic number - e.g. 35S (or S-35). Why? Q- Draw B-R diagrams for the two Li isotopes. A- The atomic # of an element doesn’t change Although the number of neutrons can vary, atoms have definite numbers of protons.

6Li 7Li 3 p+ 3 n0 2e– 1e– 3 p+ 4 n0 2e– 1e–

Practice P. 29 #1-7

Limitations to Rutherford’s Model Orbiting electrons should emit light, losing energy in the process This energy loss should cause the electrons to collapse into the nucleus However, matter is very stable, this does not happen

Bohr’s Planetary Model Electrons orbit the nucleus in energy “shells” An electron can travel indefinitely within a shell without losing energy The greater the distance between the nucleus and the shell, the greater the energy level An electron cannot exist between shells, but can move to a higher, unfilled shell if it absorbs a specific quantity of energy, or to a lower, unfilled shell if it loses energy (quantized) When all the electrons in an atom are in the lowest possible energy levels, it is in its ground state.

An atom becomes excited when one of its electrons absorb energy If enough energy is absorbed then the electron can make a quantum leap to the next energy level, if there is room When the electron returns to a lower energy state the energy is released in the form of a photon, which we see as visible light

The energy of the photon determines its wavelength or color Each element has its own frequencies of color, so it emits its own distinctive glow

Summary of Atomic Models Dalton’s “Billiard ball” model (1800-1900) Atoms are solid and indivisible. Thomson’s “Raisin bun” model (1900) Negative electrons in a positive framework. Rutherford’s “Nuclear” model (~1910) Atoms are mostly empty space. Negative electrons orbit a positive nucleus. 4) Bohr’s “Planetary” model (~1920) Negative electrons orbit a positive nucleus. Quantized energy shells 5) Quantum Mechanical model (~1930) Electron probabilities (orbitals)

Practice Pre-lab: Atomic Spectra (p. 40) p. 42 #1-3 p. 45 #6-8