Teach Me Chemistry - Lesson 4
We saw at the end of the last lesson that protons and neutrons each weigh approximately one atomic mass unit (amu, unit “u”), while electrons in comparison weigh only 1/2000th of that. Formally, an amu is defined as the exactly 1/12 the mass of carbon-12. Let’s look at how atomic number, mass number, and the numbers of the subatomic particles are related.
X represents the element’s chemical symbol found in the periodic table, while A is the atomic mass and Z is the atomic number (also sometimes known as the proton number). For instance, 
represents the element carbon - “C” in the periodic table - which has 6 protons and a mass of 12 amu.
Because we know that protons and neutrons are the only subatomic particles that count towards the atomic mass and each has a mass of one amu, we can calculate that the number of neutrons in this atom is 6.
A - Z = # neutrons
12 - 6 = 6
II. Isotopes
It is important to note that all atoms of a given element will always have the same number of protons. Carbon always has 6, oxygen always has 8, uranium always has 92, and so on. If the number of protons changed, we would have a different element entirely. Number of protons therefore gives the atom its identity.
In contrast, atoms of the same element can have different masses. For example, while most naturally occurring carbon has a mass of 12 u (carbon-12), there also exists a significant amount of carbon-13 and carbon-14, which you may have heard of in the context of radioactive dating of ancient artifacts. Atoms that have the same number of protons but different masses are called isotopes.
Where does the change in mass come from? By process of elimination - the difference can’t be in the protons because it would change the atom’s identity, and it can’t be in electrons because they don’t weigh enough - we know it must be a difference in the number of neutrons that gives us isotopes.
If we use the A minus Z method on carbon-14, we see that this molecule has 8 neutrons.
Note: it’s important not to confuse isotopes with ions! An ion, or charged atom, results from a gain or loss of electrons. In an electrically neutral atom, the number of protons will equal the number of electrons. Because electrons are negatively charged, gaining one makes the overall charge of the atom more negative, while losing one makes it positive. A sodium (Na) ion which has lost one electron would be represented this way:
Q: How many protons does this atom have?
A: Eleven (the atomic number or proton number)
Q: How many neutrons?
A: 23-11 = 12
Q: How many electrons?
A: 11-1 = 10
III. Real Atomic Masses
In practice, observed atomic masses must take into account all the masses of the naturally occurring isotopes. To calculate this value, we use a weighted average based on the relative abundances of the isotopes.
atomic mass = (fractional abundance of isotope 1 x mass of isotope 1) + (fractional abundance of isotope 2 x mass of isotope 2) + … +
(This is the same calculation you must undertake to calculate your grade point average, “weighting” each grade with the units of credit for the corresponding class.)
IV. Homework!
1. a. How many protons, neutrons, and electrons are in an atom of 
?
b. How about 
?
c. And 
?
2. If 98.892% of carbon atoms are carbon-12, with a weight of 12 u. and 1.108% are carbon-13 with a mass of 13.00335 u, calculate the atomic mass of naturally occurring carbon.