The Structure of the Atom for AP Chemistry (page 2)
Practice problems for these concepts can be found at:
- Basics Multiple Choice Review Questions for AP Chemistry
- Basics Free-Response Questions for AP Chemistry
The States of Matter
Matter can exist in one of three states: solid, liquid, or gas. A solid has both a definite shape and a definite volume. At the molecular level, the particles that make up a solid are close together and many times are locked into a very regular framework called a crystal lattice. Molecular motion exists, but it is slight.
A liquid has a definite volume but no definite shape. It conforms to the container in which it is placed. The particles are moving much more than in the solid. There are usually clumps of particles moving relatively freely among other clumps.
A gas has neither definite shape nor volume. It expands to fill the container in which it is placed. The particles move rapidly with respect to each other and act basically independently of each other.
We will indicate the state of matter that a particular substance is in by a parenthetical s, l, or g. Thus, H2O(s) would represent solid water (ice), while H2O(g) would represent gaseous water (steam).
The Structure of the Atom
The first modern atomic theory was developed by John Dalton and first presented in 1808. Dalton used the term atom (first used by Democritus) to describe the tiny, indivisible particles of an element. Dalton also thought that atoms of an element are the same and atoms of different elements are different. In 1897, J. J. Thompson discovered the existence of the first subatomic particle, the electron, by using magnetic and electric fields. In 1909, Robert Millikan measured the charge on the electron in his oil drop experiment (electron charge = –1.6022 × 10–19 coulombs), and from that he calculated the mass of the electron. Thompson developed an atomic model, the raisin pudding model, which described the atom as being a diffuse positively charged sphere with electrons scattered throughout.
Ernest Rutherford, in 1910, was investigating atomic structure by shooting positively charged alpha particles at a thin gold foil. Most of the particles passed through with no deflection, a few were slightly deflected, and every once in a while an alpha particle was deflected back towards the alpha source. Rutherford concluded from this scattering experiment that the atom was mostly empty space where the electrons were, and that there was a dense core of positive charge at the center of the atom that contained most of the atom's mass. He called that dense core the nucleus.
Our modern theory of the atom describes it as an electrically neutral sphere with a tiny nucleus at the center, which holds the positively charged protons and the neutral neutrons. The negatively charged electrons move around the nucleus in complex paths, all of which comprise the electron cloud. Table 5.1 summarizes the properties of the three fundamental subatomic particles:
Many teachers and books omit the charges on the symbols for the proton and neutron.
The amu (atomic mass unit) is commonly used for the mass of subatomic particles and atoms. An amu is the mass of a carbon-12 atom, which contains 6 protons and 6 neutrons (C-12).
Since the atom itself is neutral, the number of electrons must equal the number of protons. However, the number of neutrons in an atom may vary. Atoms of the same element (same number of protons) that have differing numbers of neutrons are called isotopes. A specific isotope of an element can be represented by the following symbolization:
X represents the element symbol taken from the periodic table. Z is the atomic number of the element, the number of protons in the nucleus. A is the mass number, the sum of the protons and neutrons. By subtracting the atomic number (p) from the mass number (p + n), the number of neutrons may be determined. For example, (U-238) contains 92 protons, 92 electrons, and (238 – 92) 146 neutrons.
Electron Shells, Subshells, and Orbitals
According to the latest atomic model, the electrons in an atom are located in various energy levels or shells that are located at different distances from the nucleus. The lower the number of the shell, the closer to the nucleus the electrons are found. Within the shells, the electrons are grouped in subshells of slightly different energies. The number associated with the shell is equal to the number of subshells found at that energy level. For example, energy level 2 (shell 2) has two subshells. The subshells are denoted by the symbols s, p, d, f, etc. and correspond to differently shaped volumes of space in which the probability of finding the electrons is high. The electrons in a particular subshell may be distributed among volumes of space of equal energies called orbitals. There is one orbital for an s subshell, three for a p, five for a d, seven for an f, etc. Only two electrons may occupy an orbital. Table 5.2 summarizes the shells, subshells, and orbitals in an atom. The chapter on Spectroscopy, Light, and Electrons, Chapter 10 has a discussion of the origin of this system.
The information above can be shown in graph form as an energy-level diagram, as shown in Figure 5.1:
Be sure to fill the lowest energy levels first (Aufbau principle) when using the diagram above. In filling orbitals having equal energy, electrons are added to the orbitals to half fill them all before any pairing occurs (Hund's rule). Sometimes it is difficult to remember the relative energy position of the orbitals. Notice that the 4s fills before the 3d. Figure 5.2 may help you remember the pattern in filling. Study the pattern and be able to reproduce it during the exam.
Following these rules, the energy-level diagram for silicon (Z = 14) can be written as shown in Figure 5.3
Although this filling pattern conveys a lot of information, it is bulky. A shorthand method for giving the same information has been developed—the electronic configuration.
The electronic configuration is a condensed way of representing the pattern of electrons in an atom. Using the Aufbau build-up pattern that was used in writing the energy-level diagram, consecutively write the number of the shell (energy level), the type of orbital (s, p, d, etc.), and then the number of electrons in that orbital shown as a superscript. For example, 1s22s1 would indicate that there are two electrons in the s-orbital in energy level (shell) l, and one electron in the s-orbital in energy level 2. Looking at the energy-level diagram for silicon above, the electronic configuration would be written as:
- silicon : 1s22s22p63s2 3p2
The sum of all the superscripts should be equal to the number of electrons in the atom (the atomic number, Z). Electronic configurations can also be written for cations and anions.
Practice problems for these concepts can be found at:
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