SAT Chemistry Atomic Structure and the Periodic Table of the Elements - Electric Nature Of Atoms

SAT Chemistry Atomic Structure and the Periodic Table of the Elements - Electric Nature Of Atoms

From around the beginning of the twentieth century, scientists have been gathering evidence about the structure of atoms and fitting the information into a model of the atomic structure.

Basic Electric Charges
The discovery of the electron as the first subatomic particle is credited to J. J. Thomson (England, 1897). He used an evacuated tube connected to a spark coil as shown in Figure 3. As the voltage across the tube was increased, a beam became visible. This was referred to as a cathode ray. Thomson found that the beam was deflected by both electrical and magnetic fields. Therefore, he concluded that cathode rays are made up of very small, negatively charged particles, which became known as electrons.


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Further experimentation led Thomson to find the ratio of the electrical charge of the elec­tron to its mass. This was a major step toward understanding the nature of the particle. He was awarded a Nobel Prize in 1906 for his accomplishment.
It was an American scientist, Robert Millikan, who in 1909 was able to measure the charge on an electron using the apparatus pictured in Figure 4.

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Oil droplets were sprayed into the chamber and, in the process, became randomly charged by gaining or losing electrons. The electric field was adjusted so that a negatively charged drop would move slowly upward in front of the grid in the telescope. Knowing the rate at which the drop was rising, the strength of the field, and the mass of the drop, Millikan was able to calculate the charge on the drop. Combining the information with the results of Thomson, he could calculate a value for the mass of a single electron. Eventually, this number was found to be 9.11 × 10-28 gram.
Ernest Rutherford (England, 1911) performed a gold foil experiment (Figure 5) that had tremendous implications for atomic structure.

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Alpha particles (helium nuclei) passed through the foil with few deflections. However, some deflections (1 per 8,000) were almost directly back toward the source. This was unex-pected and suggested an atomic model with mostly empty space between a nucleus, in which most of the mass of the atom was located and which was positively charged, and the electrons that defined the volume of the atom. After two years of studying the results, Rutherford finally came up with an explanation. He reasoned that the rebounded alpha particles must have experienced some powerful force within the atom. And he assumed this force must occupy a very small amount of space, because so few alpha particles had been deflected. He con¬cluded that the force must be a densely packed bundle of matter with a positive charge. He called this positive bundle the nucleus. He further discovered that the volume of a nucleus was very small compared with the total volume of an atom. If the nucleus were the size of a marble, then the atom would be about the size of a football field. The electrons, he suggested, surrounded the positively charged nucleus like planets around the sun, even though he could not explain their motion.
Further experiments showed that the nucleus was made up of still smaller particles called protons. Rutherford realized, however, that protons, by themselves, could not account for the entire mass of the nucleus. He predicted the existence of a new nuclear particle that would be neutral and would account for the missing mass. In 1932, James Chadwick (England) dis-covered this particle, the neutron.

Bohr Model of the Atom
In 1913, Niels Bohr (Denmark) proposed his model of the atom. This pictured the atom as having a dense, positively charged nucleus and negatively charged electrons in specific spher-ical orbits, also called energy levels or shells, around this nucleus. These energy levels are arranged concentrically around the nucleus, and each level is designated by a number: 1, 2, 3,.. .The closer to the nucleus, the less energy an electron needs in one of these levels, but it has to gain energy to go from one level to another that is farther away from the nucleus.
Because of its simplicity and general ability to explain chemical change, the Bohr model still has some usefulness today.
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Components of Atomic Structure
The chart below lists the basic particles of the atom and important information about them.
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When these components are used in the model, the protons and neutrons are shown in the nucleus. These particles are known as nucleons. The electrons are shown outside the nucleus.
The number of protons in the nucleus of an atom determines the atomic number. All atoms of the same element have the same number of protons and therefore the same atomic number; atoms of different elements have different atomic numbers. Thus, the atomic number identifies the element. An English scientist, Henry Moseley, first determined the atomic numbers of the elements through the use of X-rays.
The sum of the number of protons and the number of neutrons in the nucleus is called the mass number.
Table 1 (see page 64) summarizes the relationships just discussed. Notice that the outer­most energy level can contain no more than eight electrons. The explanation of this is given in the next section.
In some cases, different types of atoms of the same element have different masses. For example, three types of hydrogen atoms are known. The most common type of hydrogen, sometimes called protium, accounts for 99.985% of the hydrogen atoms found on Earth. The nucleus of a protium atom contains one proton only, and it has one electron moving about it. The second form of hydrogen, known as deuterium, accounts for 0.015% of Earth’s hydro­gen atoms. Each deuterium atom has a nucleus containing one proton and one neutron. The third form of hydrogen, tritium, is radioactive. It exists in very small amounts in nature, but it can be prepared artificially. Each tritium atom contains one proton, two neutrons, and on - electron.
Protium, deuterium, and tritium are isotopes of hydrogen. Isotopes are atoms of the same element that have different masses. The isotopes of a particular element all have the same number of protons and electrons but different numbers of neutrons. In all three isotopes of hydrogen, the positive charge of the single proton is balanced by the negative charge of the electron. Most elements consist of mixtures of isotopes. Tin, for example, has ten stable iso­topes, the most of any element.

The percentage of each isotope in the naturally occurring element on Earth is nearly always the same, no matter where the element is found. The percentage at which each of an element’s isotopes occurs in nature is taken into account when calculating the element’s average atomic mass. Average atomic mass is the weighted average of the atomic masses of the naturally occurring isotopes of an element.
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Calculating Average Atomic Mass
The average atomic mass of an element depends on both the mass and the relative abun­dance of each of the element’s isotopes. For example, naturally occurring copper consists of 69.17% copper-63, which has an atomic mass of 62.919 598 amu, and 30.83% copper-65, which has an atomic mass of 64.927 793 amu. The average atomic mass of copper can be calculated by multiplying the atomic mass of each isotope by its relative abundance (expressed in decimal form) and adding the results.
0.6917 × 62.919 598 amu + 0.3083 × 64.927 793 amu = 63.55 amu
Therefore, the calculated average atomic mass of naturally occurring copper is 63.55 amu. Average atomic masses of the elements listed in the Periodic Table, rounded to one decimal place for use in calculations and also in full to four decimal places, are given in the Chemical Elements table in the Tables for Reference section at the back of the book.

Valence Electrons
Each atom attempts to have its outer energy level complete and accomplishes this by bor­rowing, lending, or sharing its electrons. The electrons found in the outermost energy level are called valence electrons. The remainder of the electrons are called core electrons. The absolute number of electrons gained, lost, or borrowed is referred to as the valence of the atom.

Example: ___________________________________________
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This picture can be simplified to •Cl:, showing only the valence electrons as dots in an electron dot notation. This is called the Lewis structure of the atom. To complete its outer orbit to eight electrons, chlorine must borrow an electron from another atom. Its valence number then is 1. As stated above, when electrons are gained, we assign a - sign to this number, so the oxidation number of chlorine is -1 (more on pages 75 and 76).

Another Example: ___________________________________________
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Since sodium tends to lose this electron, its oxidation number is +1.

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