The chapter starts out with some historical perspectives, including the formulation of some fundamentally important chemical laws, then moves on the one of the foundations of modern chemistry - the atomic theory - to explain these laws. Note that this approach mirrors that of the scientific method described in chapter 1: observe nature, formulate laws to summarize the observations, then explain the laws with a theory. Make sure you understand the following laws, and how they are explained by the atomic theory: Law of Conservation of Mass, Law of Definite Proportions, and Law of Multiple Proportions.

Dalton's Atomic Theory
Make sure you know the four basic postulates of Dalton's Atomic Theory. The second postulate mentioned in the text, that all atoms of a given element are identical, is now understood to be false. Isotopes (atoms of the same element with different numbers of neutrons), which weren't known in Dalton's time, are not identical.

The text describes some ground-breaking experiments which provided us with our understanding of the structure of the atom:

1. J. J. Thomson's determination of the charge-to-mass (e/m) ratio for the electron by his cathode ray studies. This ratio is -1.76 x 10⁸ coulomb/gram. It is important to understand that Thomson didn't know either the mass of the electron or its charge individually; he only determined the ratio.
2. Robert Millikan's "oil drop" experiment, by which he determined the charge on an electron: -1.602 x 10^-18 coulomb (absolute charge); -1 (charge relative to the proton's +1 charge). Using this result with Thomson;s value above allowed the mass of the electron to be determined: 9.11 x 10^-31 kg.
3. Ernest Rutherford's experiments with alpha particle bombardment of thin metal foils. These experiments clearly provided the information which led to the formulation of the atom as a nuclear atom. Most of the atom is empty space. The vast majority of the mass of an atom, and all of the positive charge, is concentrated in a very tiny volume called the nucleus. Fig. 2.14 shows that a typical atom has a diameter of 10^-8 cm, while the diameter of a nucleus is about 100,000 times smaller: on the order of 10^-13 cm.

The atomic number (Z) is the number of protons in the nucleus. Since an atom is defined by its atomic number, all atoms of a given element have the same Z. The mass number (A) is the sum of the number of protons and neutrons. Thus, A must always be an integer. Don't confuse the mass number with the atomic weight, which will be discussed in chapter 3. Z is usually appended as a lower-left subscript to the atomic symbol and N is an upper-left superscript. Since the atomic symbol and the atomic number are redundant, a notation which doesn't require the use of subscripts and superscripts is something like Cl-35, where the 35 represents the mass number. The number of neutrons can be obtained by subtracting the atomic number from the mass number. Thus, since the atomic number of chlorine is 17, there are 35-17=18 neutrons in a Cl-35 atom.

Molecules and Ions
This section of the text discusses various types of formulas for compounds and several ways to depict the structure (space-filling models, ball-and-stick models, etc.). These ideas should be a review of what was covered in the Organic Chemistry supplement.

Although the difference between covalent and ionic bonds will be covered in detail in chapter 8 , you should know that only covalent bonds result in molecules - a tightly-bound group of atoms which can exist as a discrete unit. Ionic compounds do not contain molecules - they consist of cations (positively-charged atoms) and anions (negatively-charged atoms) which exist in a three-dimensional arrangement called a crystal lattice (this will be discussed more fully in chapter 10 .) The unit making up an ionic compound is called a formula unit. Thus, a chemist talks about a molecule of the covalent substance water (H₂O), but a formula unit of the ionic substance sodium chloride (NaCl).

Introduction to the Periodic Table
Although the periodic table will be discussed in much more detail in chapter 7 , there are some basic ideas which need to be presented in chapter 2 because some of the terminology will be used throughout the next several chapters.

The black staircase line in figure 2.21 separates the nonmetals (to the right of the staircase line) from the metals (to the left of the line.) Many elements which are along the line exhibit properties of metals and nonmetals. They are called metalloids (or, sometimes, semi-metals.)

The elements in the periodic table are arranged vertically in groups (or families) -- Groups IA, IIA, etc.-- and horizontally in periods (or rows). Some of the groups have common names: Group IA is known as the alkali metals, Group IIA elements are alkaline earth metals, Group VIIA is the halogens, and Group VIIIA are the noble gases.

Naming Compounds
The section on naming compounds is very complete. It will not be reproduced here. Make sure you know all of the names (by heart!!) in Tables 2.3 through 2.8. The flow charts are very useful in helping you to learn what decisions need to be made before the appropriate name is assigned. Study these carefully.

Last modified June 16, 1997