The teacher should provide students with picture of a general and simplified structure of the atom, showing the nucleus composed of protons and neutrons, and electrons outside the nucleus. Protons are positively charged and are held together in the nucleus by the neutrons (proton glue) which has no charge. Electrons are negatively charged and exist outside the nucleus. Both protons and electrons have equal but opposite charge magnitude. Protons and neutrons have relative masses of ~1.00g and electron has a relative mass of 0.00 g thus the mass of atoms is roughly equivalent to the combined mass of the protons and neutrons. (See Table 1 for some general information on protons, neutrons, and electrons).
Exercise Determining the number Particles in an Atom
The periodic table can be used to determine the number of protons, electrons, and neutrons in a particular element. For any given element the number of protons (p) is equal to the atomic number (z) of the element (p = z) and the number of electrons (e) equals the atomic number minus the charge (e = z – c). The number of neutrons (n) is equal to the mass number (m) minus the atomic number (n = m – z).
In this exercise students must complete a table by calculating the number of protons, neutrons, and electrons (See worksheet 2).
Energy and the Atom
While the general structure of the atom was being discovered, physicists were performing experiment that would revolutionize how matter is viewed. They observed that each element, when heated, produced a characteristic color with discrete wavelengths of light (line spectrum) that was unlike the uniform rainbow spectrum of white light passing through a prism. With the exception of hydrogen, the line spectra of most elements are very complex. Consequently, Niels Bohr used the hydrogen atom to explain a model that considered line spectrum. The solar system model, as it is called, assumes that electron move around the nucleus in certain allowed orbits (which are also called energy levels or shells). Electrons thus retain definite energy characteristic of the orbit in which it is moving. The further a shell is from the nucleus, the more energy an electron will possess. Electron can move from the lowest possible Energy State (E1; ground state) when they absorb energy (i.e. when heated) to a higher energy state (E2; excited state). When the electron returns to the ground state, it emits a definite amount of energy (DðE) equal to the difference in energy between the ground and excited states (DðE = E2 - E1). The characteristic color or wavelength is explained using the equation, DðE = h × c (lð, where h and c are constants and lð is the wavelength. Since each element is capable of emitting a unique amount of energy, each will give a unique wavelength.
The actual model is more complex than the Bohr’s solar system model. Teachers can find additional information on the wave-mechanical model most introductory chemistry textbooks, or see Mortimer’s Chemistry listed in the reference section. It will suffice, however, to simply discuss the Bohr’s model of the hydrogen atom so that student have a general idea of orbits, shells, atomic energy, etc.
This section of the unit will introduce students to the Bohr’s model of the hydrogen atom and help students develop an appreciation for electron orbits and the related energy. It will focus mainly on a lecture and a laboratory experiment.
This section of the unit will partially fulfill Content Standards on scientific inquiry of New Haven Public Schools Academic Performance Standard by:
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1. Designing and conducting scientific investigation while taking proper safety precaution into consideration
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2. Using mathematics in scientific inquiry
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3. Recognizing that the results of scientific inquirynew knowledge and methodsemerge from different type of investigations and public communication among scientist.
Instructional Technique
The time limit for this section is 3-4 45 minutes class periods. It centers on a lecture/class discussion and a laboratory exercise.
Lecture on the Bohr’s Model for the Hydrogen Atom
The important concepts to get across are described in the section summary above. To begin discussing the solar system model it will be helpful to show a figure of the solar system to be used as an analogy for explaining the hydrogen atom. Student should be able to visualize the planets as electrons moving in specific orbits and the sun as the nucleus. They must also understand why each element will produce a unique line spectrum. This is accomplished by describing each element as a solar system with varying number of planets at unique distances from their sun. If each unit distance from the sun is a specific energy value, when students calculate DðE for various excitement/relaxation events they will observe unique series of wavelengths for each solar system.
Flame Test Lab Exercise
The laboratory exercise to be used can be obtained in Prentice Hall Chemistry: The Study of Matter, page 127. In this exercise students will observe unique color emission for different metallic ions (Na+, K+, Li+, Ca2+, Sr2+, Ba2+, Cu2+) when heated and, base on this observation, must determine the identity of an unknown substance.
Teacher General Background on Electron Structure of the Atom
The rotation of electrons around the nucleus is much more complex than the model proposed by Bohr. All electrons do not move in circular pattern, as do planets but instead move in spherical and dumb-bell shaped spatial regions around the nucleus that are called orbitals. There are 16 different orbitals (1 s orbital; 3 p orbitals -px, py, pz; 5 d orbitals -dyz, dxz, dxy, dx2-y2, d z2, and 7 different f orbitals). Orbitals exist in subshells that in turn exist in shells (these terms are explained later). Unlike the planetary orbits where the exact location of a planet can be pin-point at any given time, electron orbitals represent a region of space around the nucleus where there exists a high density of electron. It gives the probability of an electron be at a specific point in space at a given time. Each orbital has a distinct shape and orientation in space. s orbitals are spherical around the nucleus while the other three are dumb-bell shaped (for a pictorial representation of each see Jespersen Chemistry, page 15-16). As stated previously, the number of each type of orbital may vary. There is 1 s, 3 p, 5 d, and 7 f orbitals, each aligned with the x-, y-, and z-axis differently (see Jespersen page 15-16). Each orbital holds a maximum of two electron, thus s orbitals hold a total of 1 ( 2 = 2 , p holds 3 ( 2 = 6, d holds 5 ( 2 = 10, and f hold 7 ( 2 = 14 electrons.
The modern model of the atom has a nucleus of positively charged protons surrounded by orbitals of negatively charged electrons. The electrons, moving in their orbitals, are distributed at distinct Principal Energy Levels (or shells). These shells are distinguished by their Principal Quantum Numbers (n) 1 for the first shell (closest to the nucleus), 2 for the second, 3 for the third, and so on. Each shell can have a maximum of n2 orbitals (i.e. the second shells can have 22 or 4 orbitals) and 2n2 electrons.
Energy Sublevels or subshells exists within shells and are the direct vessels for orbitals (orbitals in subshells; subshells in shells) and are represented by their Azimuthal Quantum Number (l). l begins with 0 and can move up incrementally by 1 but can never be greater than n-1, thus there can be a maximum of n subshells. There are four different types of subshells that are of importance to high school chemistry; these are s, p, d, and f subshells with l = 0, 1, 2, and 3, respectively.
Magnetic Quantum Numbers (ml) are used to represent the orbitals. ml has a value from - l ...o...+ l . For example, for l =1 (p subshells) ml =-1,0,1 where these numbers represent px, pz, and py, respectively. Since each orbital can carry a maximum of 2 electrons there must be a way to distinguish the two electrons. Pair electrons must have opposite spins that are given a Spin Quantum Number2 (ms) of +½ or –½.
