WebLecture I: Atoms and Isotopes
WebLecture II: A History of Chemistry before the Curies (Optional: some of this information has also been incorporated into other weblectures)
Kotz and Triechel, Chemistry and Chemical Reactivity Chapter 2: Sections 1-4.
- 2.1 Atomic components When Dalton resurrected the ancient Greek concept of atomi to explain whole-number ratios of elements in compounds, and fixed-formula compound structure, he assumed, as Democritus had, that the elemental atoms were indivisible, and this "rule" persisted until observations of radiation and electricity convinced the Curies and Thompson that atoms themselves contained components that could be rearranged or ejected from inside. Rutherford's alpha-particle experiment offered further evidence that atoms are mostly space, and that the component stable particles (protons, neutrons, and electrons) are very small compared to the overall size of the atom.
- 2.2 Atomic number and mass Atomic number identifies the element, atomic mass identifies (with atomic number) the isotope. Atomic number is simply the number of protons in the atom. All carbon atoms, for instance, have 6 protons, regardless of the number of electrons or neutrons they may also have at any given instant. It is the number of protons that makes the atom a particular element. Atomic mass is calculated using atomic mass units, where an amu = 1/12 * the mass of a carbon-12 atom with six protons and six neutrons. Because the number of neutrons in atoms of the same element may vary, the average atomic mass of an sample can be a fractional value. We can have a carbon sample, where some atoms have atomic mass 12 (6 protons + 6 neutrons), and other carbon atoms with atomic mass 14 (6 protons + 8 neutrons) so the [total mass]/[number of atoms] will be around 12.11 amu. Number of protons and total number of "nucleons" (protons plus neutrons) identify isotopes.
- 2.3 Isotopes Isotopes are atoms of the same element (same number of protons) with different numbers of neutrons. Because the number of neutrons differs, the mass differs. When mass is given in amu, the mass for a single isotope will be a whole number. In any naturally-occurring sample of a element with multiple isotopes, there will be some atoms of each. You need to be able to figure the relative abundances of different isotopes given the atomic number, the atomic weight, and the number of neutrons in each isotope.
- 2.4 The atomic weight is an average based on experimental data: the masses of all the atoms of all the isotopes divided by the total number of atoms. You may discover differences in atomic weights or atomic mass if you look at periodic tables created decades apart. One of the reasons is that our definition of an atomic mass unit has changed from a basis on the mass of hydrogen to a basis as 1/12 of the mass of carbon-12. While this can be confusing, the key to chemistry is proportionality -- which doesn't change when units change. We use atomic weights when we attempt to analyze or predict the presence of an element in a sample.
Homework problems: Please visit the Moodle for the current assignment and posting instructions.
Lab: Equipment and Chemicals
Please read through the chapters on equipping your home lab and safely storing chemicals in Illustrated Guide to Hom Chemistry Experiments, chapters 3 and 4. We will discuss lab safety in the AP/lab chat. You should be acquiring the chemicals for this year's labs and be ready to start developing and demonstrating lab skills and safe lab practices.
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