Unit 3 - The Atom / Radioactive Decay

Introduction:

Unit 3 switches gears entirely and focuses on the early ideas and development of the atom as discussed in chapter 3. We will discuss several men from ancient times to modern times and their contributions and the major focus of the chapter will be on the composition of the atom itself and the experiments that led to the developments. We will also talk about the internal makeup of the atom and briefly touch on some principles of nuclear chemistry from chapter 21.

Student Objectives & Assignments – Unit 3

3.1

a. The idea of atoms has been around since the time of the ancient Greeks. In the nineteenth century, John Dalton proposed a scientific theory of atoms that can still be used to explain properties of most chemicals today.

b. Matter and its mass cannot be created or destroyed in chemical reactions.

c. The mass ratios of the elements that make up a given compound are always the same, regardless of how much of the compound there is or how it was formed.

d. If two or more different compounds are composed of the same two elements, then the ratio of the masses of the second element combined with a certain mass of the first element can be expressed as a ratio of small whole numbers.

3.2

a. Cathode-ray tubes supplied evidence of the existence of electrons, which are negatively charged subatomic particles that have relatively little mass.

b. Rutherford found evidence for the existence of the atomic nucleus by bombarding gold foil with a beam of positively charged particles.

c. Atomic nuclei are composed of protons, which have an electric charge of +1, and (in all but one case), neutrons, which have no electric charge.

d. Atomic nuclei have radii of about 0.001 pm (pm=picometers; 1 pm x 10-12 m), and atoms have radii of about 40-270 pm.

3.3

a. The atomic number of an element is equal to the number of protons of an atom of that element.

b. The mass number is equal to the total number of protons and neutrons that make up the nucleus of an atom of that element.

c. The relative atomic mass unit (amu) is based on the carbon-12 atom and is a convenient unit for measuring the mass of atoms. It equals 1.660540 x 10-24 g.

d. The average atomic mass of an element is found by calculating the weighted average of the atomic masses of the naturally occurring isotopes of the element.

e. Avogadro’s number is equal to approximately 6.022 x 1023. A sample that contains a number of particles equal to Avogadro’s number contains a mole of those particles.

21.2

a. Radioactive nuclides become more stable by radioactive decay.

b. Alpha, beta, positron and gamma emission and electron capture are types of radioactive decay. The type of decay is related to the nucleon content and the energy level of the nucleus.

Ohio Department of Education - Chemistry Standards:

The physical science syllabus included properties and locations of protons, neutrons and electrons, atomic number, mass number, cations and anions, isotopes and the strong nuclear force that hold the nucleus together. In this course, the historical development of the atom and the positions of electrons are explored in more detail.

Atomic models are constructed to explain experimental evidence and make predictions. The changes in the atomic model over time exemplify how scientific knowledge changes as new evidence emerges and how technological advancements like electricity extend the boundaries of scientific knowledge. Thompson’s study of electrical discharges in cathode-ray tubes led to the discovery of the electron and the development of the plum pudding model of the atom. Rutherford’s experiment, in which he bombarded gold foil with α-particles, led to the discovery that most of the atom consists of emptyspace with a relatively small, positively charged nucleus. Bohr used data from atomic spectra to propose a planetary model of the atom in which electrons orbit the nucleus, like planets around the sun. Later, Schrödinger used the idea that electrons travel in waves to develop a model in which electrons travel randomly in regions of space called orbitals (quantum mechanical model).

Based on the quantum mechanical model, it is not possible to predict exactly where electrons are located but there is a region of space surrounding the nucleus in which there is a high probability of finding an electron (electron cloud or orbital). Data from atomic spectra (emission and absorption) gives evidence that electrons can only exist at certain discrete energy levels and not at energies between these levels. Atoms are usually in the ground state where the electrons occupy orbitals with the lowest available energy. However, the atom can become excited when the electrons absorb a photon with the precise amount of energy (indicated by the frequency of the photon) to move to an orbital with higher energy. Any photon without this precise amount of energy will be ignored by the electron. The atom exists in the excited state for a very short amount of time. When an electron drops back down to the lower energy level, it emits a photon that has energy equal to the energy difference between the levels. The amount of energy is indicated by the frequency of the light that is given off and can be measured. Each element has a unique emission and absorption spectrum due to its unique electron configuration and specific electron energy jumps that are possible for that element. Being aware of the quantum mechanical model as the currently accepted model for the atom is important for science literacy as it explains and predicts subatomic interactions, but details should be reserved for more advanced study.

Electron energy levels consist of sublevels (s, p, d and f), each with a characteristic number and shape of orbitals. The shapes of d and f orbitals will not be assessed in high school. Orbital diagrams and electron configurations can be constructed to show the location of the electrons in an atom using established rules. However, the names of these rules will not be assessed. Valence electrons are responsible for most of the chemical properties of elements. In this course, electron configurations (extended and noble gas notation) and orbital diagrams can be shown for any element in the first three periods.

Although the quantum mechanical model of the atom explains the most experimental evidence, other models can still be helpful. Thinking of atoms as indivisible spheres is useful in explaining many physical properties of substances, such as the state (solid, liquid or gas) of a substance at room temperature. Bohr’s planetary model is useful to explain and predict periodic trends in the properties of elements.

Note: Quantum numbers and equations of de Broglie, Schrödinger and Plank are beyond the scope of this course.

Nuclear Reactions:

The basics of nuclear forces, isotopes, radioactive decay, fission and fusion were addressed in the physical science syllabus. In chemistry, specific types of radioactive decay and using nuclear reactions as a source of energy are addressed. Radioactive decay can result in the release of different types of radiation (alpha, beta, gamma, positron) each with a characteristic mass, charge and potential to ionize and penetrate the material it strikes. Beta decay results from the decay of a neutron and positron decay results from the decay of a proton.

When a radioisotope undergoes alpha, beta or positron decay, the resulting nucleus can be predicted and the balanced nuclear equation can be written.

Nuclear reactions, such as fission and fusion, are accompanied by large energy changes that are

much greater than those that accompanychemical reactions. These nuclear reactions can

theoretically be used as a controlled source of energy in a nuclear power plant. There are

advantages and disadvantages of generating electricity from fission and fusion

World of Chemistry Video for Class - Click on #6 "The Atom" - make sure to disable your popup blocker.

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