In the previous lesson, we learned about the subatomic particles: protons, neutrons, and electrons. In this lesson, we will try to get a more intimate understanding of these little fellas and see how the periodic table can help us determine these numbers for an element at a glance!
As you might recall, protons and electrons are positively and negatively charged, respectively. Neutrons, as their name implies, are electrically neutral; they have no charge. Although we like to think that the electron is completely massless, this is not entirely true. A single electron has a mass of 9.02 x 10-28 g, while the proton and neutron have roughly equal masses (neutrons are slightly more massive) at 1.67 x 10-24 g.
These numbers are not useful to illustrate the difference in size. Instead, we can say that if we assign a mass of 1 to the electron, then the proton and neutron would be 1836 TIMES greater in mass! That’s like comparing a pineapple to a car!
We call the number of protons within an atom’s nucleus—which is always unique to an atom of a given element—its atomic number, which is denoted by the letter Z. The nuclear mass, or the mass of the particles in the nucleus (not counting the electrons), is called an atom’s mass number and is denoted by the letter A. We can, therefore, determine the number of neutrons in an atom by subtracting the atomic number from the mass number.
There are three ways we can represent an atom, shown below. You might notice that none of these representations show the number of electrons. This is because atoms are electrically neutral, meaning that the number of protons, or the atomic number, is equal to the number of electrons.
Another way we can represent atomic nuclei is by using Bohr-Rutherford diagrams, which are graphical representations of the insights provided by the two researchers to atomic theory. The diagram consists of concentric circles representing energy levels, typically surrounding some representation of the nuclear contents: the protons (p+) and neutrons (n0). Below are the Bohr-Rutherford representations for a few atoms. Per Bohr’s model, the first energy level can only accommodate 2 electrons; the second, 8; and the third, 18.
As you might realize from the masses of subatomic particles, individual atoms are very small! Therefore, it’s not practical to represent their masses using relatively gigantic units such as the gram, as we would be dealing with extremely small decimals. Instead, we define the masses of elements relatively, that is, compared to something else. The atomic mass unit (u), is defined as 1/12 the mass of a carbon-12 atom. Therefore, neon, which has mass ~20 on the periodic table, is 20/12 the mass of a carbon-12 atom. Although we know through experimentation that one u = 1.66 x 10-27 kg, for calculation purposes, we can always use atomic mass units.
You might notice that atomic masses in the periodic table are decimals, despite being relatively close in magnitude to the mass number. Now, if electrons are negligible in mass, then how come the mass number is not the same as the element’s atomic mass? The answer to that question lies in isotopes. Originally proposed by Frederick Soddy, a colleague of Ernest Rutherford, isotopes are atoms of an element that have the same number of protons but a different number of neutrons. The masses we see on the periodic table are, therefore a weighted average of the masses of the isotopes and their respective abundance. Isotopes have similar chemical properties due to having the same number of protons and neutrons; their major difference is in mass.
Hydrogen, for example, has three different isotopes: Hydrogen-1, Hydrogen-2, and Hydrogen-3. Usually, isotopes do not have unique names as they share the same atomic number. Hydrogen is somewhat exceptional in that its H-2 isotope is called deuterium (D), and the H-3 isotope is called tritium (T). The atomic mass of hydrogen, however, is ~1.008 u, much closer to H-1 than to H-2 or H-3. This might hint you towards the idea that isotopes are not present equally in nature; this is true, the relative presence of a given isotope in nature is called its isotopic abundance. Over 99% of naturally occurring hydrogen is H-1, so it’s no wonder that the mass is so close to one.
Scientists use a device called a mass spectrometer to determine the isotopes present in a natural sample of an element. Its function is shown in the diagram below:
Some isotopes have unstable nuclei (not all, though! Oxygen has three stable isotopes: O-16, O-17, and O-18). This nuclear instability leads to spontaneous nuclear decay, called radioactivity. These unstable isotopes are specifically called radioisotopes.
When they decay, these radioisotopes give off three types of nuclear radiation—some very dangerous to living cells, others less so. In general, chemical reactions proceed towards more stable products. Radioactive decay is no different; it converts an unstable nucleus into a more stable nucleus by giving off radiation.
Something important to note when solving problems involving these types of radiation is that the mass numbers (A) and atomic numbers (Z) must always be equal on both sides of the equation!
The rate of radioactive decay varies among isotopes and is quantified by their half-life. Half-life refers to the amount of time it takes for 1/2 of all radioactive atoms to decay into more stable products. Carbon-14 is a well-known radioisotope. It occurs naturally in the atmosphere and can react with oxygen to form carbon dioxide. Humans cannot produce carbon-14 on their own; they can only gain it by consuming other life. Plants, on the other hand, can absorb carbon dioxide, storing it as glucose through photosynthesis.
When a living organism dies, it stops intaking material, so there is no more influx of carbon-14. Given that carbon-14 is a radioisotope, it begins to decay, typically into carbon-12. The half-life of carbon-14 is 5730 years. We can, therefore, approximate the time since a living organism has died using carbon-14 dating, which compares the ratio of C-14 / C-12, which is ~1:1 at the time of death but gradually decreases over time.
Finally, although they often aren’t covered in grade 11, it’s good to know about nuclear fusion and fission. These are two specific types of radioactive decay, and the differences between them are detailed as follows: