Chapter 5 – Summary
5.1 Early Atomic Theory: Dalton’s Model of the Atom
The ancient Greeks proposed that matter consists of extremely small particles called atoms. Dalton postulated that matter is composed of exceedingly small, indivisible particles called atoms. He also stated an atom is the smallest unit of an element that can participate in a chemical change and that each element has a characteristic type of atom that differs in properties from atoms of all other elements, and that atoms of different elements can combine in fixed, small, whole-number ratios to form compounds. Samples of a particular compound all have the same elemental proportions by mass. When two elements form different compounds, a given mass of one element will combine with masses of the other element in a small, whole-number ratio. During any chemical change, atoms are neither created nor destroyed.
5.2 Electric Charge
Ancient Greek philosopher, Thales of Miletus (624–546 BCE), recorded that when amber (a hard, translucent, fossilized resin from extinct trees) was vigorously rubbed with a piece of fur, a force was created that caused the fur and the amber to be attracted to each other. The rubbed amber would not only attract the fur, and the fur attract the amber, but they both could affect other (nonmetallic) objects, even if not in contact with those objects.
The English physicist William Gilbert (1544–1603) completed similar experiments with similar findings. It was concluded there were two types of electric charge – positive and negative. These forces are repulsive when the same type of charge exists on two interacting objects. These forces are attractive when the charges are of opposite types. The force acts by contact or induction. The magnitude of the force decreases (rapidly) as the objects move further away from each other. Whereas, the magnitude of the force increases (rapidly) as the objects move closer to each other. Finally, not all matter is affected by this electric force.
5.3 Subatomic Particles of the Atom
Although no one has actually seen the inside of an atom, experiments have demonstrated much about atomic structure. Thomson’s cathode ray tube showed that atoms contain small, negatively charged particles called electrons embedded in a positive atomic space. Millikan discovered that there is a fundamental electric charge—the charge of an electron. Rutherford’s gold foil experiment showed that atoms have a small, dense, positively charged nucleus; the positively charged particles within the nucleus are called protons. Chadwick discovered that the nucleus also contains neutral particles called neutrons.
5.4 Defining the Nuclear Atom
Experimental observations provided insight into the structure of the nuclear atom. The majority of the atom’s structure is made up of empty space, with a centrally located, very concentrated nucleus. The nucleus contains positively charged protons and neutrally charged neutrons. Combined, these two subatomic particles account for the majority of the mass in a given atom. The negative electrons, which contribute very little to the overall mass of the atom, are in orbit around the nucleus within the empty space. The type of element is determined by the number of protons in the nucleus of the atom. Every element has a specific atomic number that equals the number of protons it contains; an element’s atomic number can sourced from the periodic table.
The number of protons, neutrons, and electrons an atom contains differentiates one type of atom from the next. In any given atom of an element when the number of positively charged protons and the number of negatively charged electrons are the same, the atom is neutral. An atom that loses or gains electrons is called an ion. An atom that gains one or more electrons will exhibit a negative charge and is called an anion. Positively charged atoms called cations are formed when an atom loses one or more electrons.
5.5 Isotopes of the Elements
Isotopes are various forms of the same element that have the same number of protons but a different number of neutrons. A neutral atom must contain the same number of positive and negative charges, so the number of protons equals the number of electrons. Therefore, the atomic number not only indicates the number of protons but also the number of electrons in an atom. The total number of protons and neutrons in an atom is called its mass number (A). Atomic symbols for elements are used to represent specific isotopes of atoms and include their mass number (A) in superscript, atomic number (Z) in subscript, followed by the element symbol.
In summary:
[latex]\begin{array}{r @ {{}={}} l} \text{atomic number (Z)} & \text{number of protons} \\[1em] \text{mass number (A)} & \text{number of protons + number of neutrons} \\[1em] \text{A - Z} & \text{number of neutrons} \end{array}[/latex]
When reading a specific isotope symbol, it is read as “element, mass number.” For instance, in the case of magnesium, 24Mg is read as “magnesium 24,” and can be written as “magnesium-24” or “Mg-24.”
5.6 Atomic Mass
An atom consists of a small, positively charged nucleus surrounded by electrons. The nucleus contains protons and neutrons; its diameter is about 100,000 times smaller than that of the atom. The mass of one atom is usually expressed in atomic mass units (amu), which is referred to as the atomic mass. An amu is defined as exactly [latex]\frac{1}{12}[/latex] of the mass of a carbon-12 atom and is equal to 1.6605 × 10−24 g.
Protons are relatively heavy particles with a charge of 1+ and a mass of 1.0073 amu. Neutrons are relatively heavy particles with no charge and a mass of 1.0087 amu. Electrons are light particles with a charge of 1− and a mass of 0.00055 amu. The number of protons in the nucleus is called the atomic number (Z) and is the property that defines an atom’s elemental identity. The sum of the numbers of protons and neutrons in the nucleus is called the mass number and, expressed in amu, is approximately equal to the mass of the atom.
Isotopes of an element are atoms with the same atomic number but different mass numbers; isotopes of an element, therefore, differ from each other only in the number of neutrons within the nucleus, so each isotope will have a slightly different atomic mass. When a naturally occurring element is composed of several isotopes, the atomic mass of the element represents the average of the masses of the isotopes involved, and this number is represented on a periodic table for each element. Average atomic masses and percent abundance for each element’s isotopes can be calculated using formulas.
Attribution & References
Except where otherwise noted, this page is adapted by Jackie MacDonald from “2. 1 Early ideas in atomic theory”, “2.2 Evolution of Atomic Theory” and “2.3 Atomic Structure and Symbolism” In Chemistry 2e (OpenStax) by Paul Flowers, Klaus Theopold, Richard Langley & William R. Robinson is licensed under CC BY 4.0. Access for free at Chemistry 2e (OpenStax)
- Electric Charge (5.2) is adapted from “Electric charge” In University Physics Volume 2 (Open Stax) by Samuel J. Ling, William Moebs, Jeff Sanny is licensed under CC BY 4.0. Access for free at University Physics Volume 2 (Open Stax)
- Defining the Nuclear Atom (5.4) is adapted from “The structure of the atom ” In Astronomy 2e (Open Stax) by Andrew Fraknoi, David Morrison, Sidney Wolff is licensed under CC BY 4.0. Access for free at Astronomy 2e (Open Stax)