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Atoms: The Building Blocks of Matter. Chapter 3. Foundations of the Atomic Theory.
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Atoms: The Building Blocks of Matter Chapter 3
If you crushed a sugar cube, you would find a number of small fragments that were still sugar. If you crushed a fragment, you would find a number of small particles that were still sugar. How long could you divide the sugar until the particles were no longer sugar?
This was the approach that Democritus took in explaining the atom. He said that you could take a pair of shears and cut a piece of copper in two and sometime, you would reach a piece that couldn’t be cut anymore. He named this particle an atom meaning indivisible.
In the late 1700’s, chemists had named a number of elements. To them, an element was a substance that could not be broken down into simpler substances by ordinary chemical means. With this understanding, thousands of experiments, and better technology, scientists discovered several laws about chemical reactions.
The Law of Conservation of Mass states that matter is neither created nor destroyed in ordinary chemical or physical reactions.
The Law of Definite Proportions states that a compound always contains the same elements in the same proportions by mass.
The Law of Multiple Proportions states that when two or more of the same elements make up several different compounds, the ratios of the reacting elements always compare in the ratios of small whole numbers.
Dalton’s Atomic TheoryIn 1808, an English chemistry teacher named John Dalton proposed the atomic theory. From data gathered in his student’s experiments, he explained the theories mentioned above and laid the foundation for understanding the atom.
Dalton’s Atomic Theory can be summed up in the following statements: • All matter is composed of small particles called atoms. • Atoms of the same element are identical in size, mass, and other properties; atoms of different elements differ in size, mass, and properties. • Atoms cannot be subdivided, created, or destroyed. • Atoms of different elements combine in small number ratios to form compounds. • In chemical reactions, atoms are combined, separated, or rearranged.
Not all aspects of Dalton’s theory have proven to be correct. We now know that atoms can be subdivided and that atoms of the same element can have different masses. However the most important concept that all matter is made of elements and that different elements have different properties remains unchanged.
As the atom was studied in depth, it became clear that it contained sub-particles. So today, we define the atom as the smallest part of an element that retains the chemical properties of that element.
All atoms have two general regions, a tiny, massive, positively charged nucleus surrounded by a negatively charged electron cloud. The nucleus is made up of positively charged protons and neutrons that have no charge.
Gases at atmospheric pressure do not conduct electricity well, however, gases at very low pressures do conduct. A glowing current will pass from a negatively charged cathode to a positively charged anode. This stream was called a cathode ray and the device was called a cathode ray tube.
Experiments using the tube showed that an object placed in the tube cast a shadow and the cathode ray could cause a paddle wheel to roll along rails through the tube. This indicated that the ray was of a particle nature. Cathode rays are deflected by a magnetic field and the rays were deflected away from a negatively charged object. This would indicated that they carry a negative charge.
In 1897, Joseph John Thomson measured the ratio of the charge of the cathode particles to their mass and found it always to be the same. Thomson concluded that cathode rays where made up of identical, negatively charged particles which were later named electrons.
In 1909, Robert Millikan performed the famous oil drop experiment to establish the mass of an electron to be 1/2000 of a hydrogen atom which was the smallest particle known. Since, this has been established at 9.109 x 10(-31) kg or 1/1837 of a hydrogen atom. By the same token, Millikan determined the charge of an electron.
Based on the discoveries made about electrons, scientists made two inferences: • Because atoms are electrically neutral, they must contain a positive charge to balance the electrons. • Atoms must contain other particles that account for most of their mass.
In 1911, New Zealander Earnest Rutherford and his associates Hans Geiger and Ernest Marsden bombarded a thin gold foil with alpha particles emitted from a radioactive source. They expected an evenly charged force field in the foil so they planned that the particles would pass straight through. When the detector was studied, they were greatly surprised to find about 1 in 8000 particles bounced straight back! Rutherford exclaimed that this would be like firing a 15 inch artillery shell into a piece of tissue paper and have it bounce back.
Rutherford concluded that atoms must contain a very small, dense, positively charged nucleus. The nucleus of the atom is very small. If it were the size of a marble, the atom would be larger than a football field.
Except for the simplest type of hydrogen, the nucleus of atoms contains two types of particles: • They contain positively charged protons. The number of protons identifies the atom. • They also contain neutrons which are equal in mass to protons but do not carry a charge.
Since most atoms in nature carry no net charge, the number of protons and electrons in a resting atom is the same. Otherwise, it would be shocking!
Since particles with the same charge repel each other, we would expect the protons crowded together in the nucleus to be unstable. However, when the protons come into extremely close range there is a strong attraction between them. These strong, short range, proton- proton, neutron-proton, and neutron-neutron forces are referred to as nuclear forces.
The radius of the atom is measured from the center of the nucleus to the outer edge of the electron cloud which surrounds the nucleus. This distance varies from 40 to 270 pm. Nuclei have extremely high densities. In fact, the density of a nucleus is about 200,000,000 metric tons per cubic centimeter.
Counting Atoms Section Three
The atomic number of an element is the number of protons in the nucleus. This number identifies the element. It will be found near the top of each grid on the periodic table. For instance, hydrogen is 1, lithium is 3, carbon is 6, and silver is 47.
The relative atomic mass of an element is the mass of an atom of that element as compared to an atom of carbon-12. Since most elements have several isotopes, this becomes a weighted average (average atomic mass or atomic mass) and may be followed by several decimal places.
The mass number is the number of protons and neutrons in the nucleus. It is derived by rounding off the atomic mass.
Naturally occurring elements may contain variations that have differing numbers of neutrons. Since they have the same number of protons, they are still the same element but they have a different mass number. These variations are called isotopes.
Most of the hydrogen in the world is protium or hydrogen-1. This isotope of hydrogen has one proton in the nucleus surrounded by one electron. A small bit of hydrogen, however, has a nucleus containing one proton and one neutron. This hydrogen-2 is called deuterium. Some radioactive hydrogen or tritium (hydrogen-3) has one proton and two neutrons in the nucleus.
The mole is the SI unit for measuring the amount of substance. One mole is precisely the amount of substance that contains the number of atoms in 12 grams of Carbon-12.
This number of atoms is called Avogadro’s Number which is 6.022137 x 1023 atoms.
The molar mass of an element is equal to the atomic mass stated in grams.
It is important to understand conversion mole quantities to work efficiently in chemistry.
