Chemistry Essentials for Dummies
Chapter 2. What's In an Atom?
Isotopes and Ions
The number of protons in an atom determines which element you have. But sometimes the number of neutrons or electrons varies, so you see several different versions of the atoms of that element. In this section, I introduce you to two variations: isotopes and ions.
Isotopes: Varying neutrons
REMEMBER. The atoms of a particular element can have an identical number of protons and electrons but varying numbers of neutrons. If they have different numbers of neutrons, then the atoms are called isotopes.
Hydrogen is a common element here on Earth. Hydrogen’s atomic number is 1 — its nucleus contains 1 proton. The hydrogen atom also has 1 electron. Because it has the same number of protons as electrons, the hydrogen atom is neutral (the positive and negative charges have canceled each other out).
Most of the hydrogen atoms on Earth contain no neutrons. You can use the symbolization in Figure 2-1 to represent hydrogen atoms that don’t contain neutrons, as shown Figure 2-7a shows.
However, approximately one hydrogen atom out of 6,000 contains a neutron in its nucleus. These atoms are still hydrogen, because they each have one proton; they simply have a neutron as well, which most hydrogen atoms lack. So these atoms are called isotopes. Figure 2-7b shows an isotope of hydrogen, commonly called deuterium. It’s still hydrogen, because it contains only one proton, but it’s different from the hydrogen in Figure 2-7a, because it also has one neutron. Because it contains one proton and one neutron, its mass number is 2 amu.
There’s even an isotope of hydrogen containing two neutrons. This one’s called tritium, and it’s represented in Figure 2-7c. Tritium is extremely rare, but it can easily be created.
Figure 2-7 also shows an alternative way of representing isotopes: Write the element symbol, a dash, and then the mass number.
Figure 2-7: The isotopes of hydrogen.
REMEMBER. Now you may be wondering, “If I’m doing a calculation involving atomic mass, which isotope do I use?” Well, you use an average of all the naturally occurring isotopes of that element — but not a simple average. Instead, you use a weighted average, which takes into consideration the abundances of the naturally occurring isotopes. You find this number on the periodic table.
For hydrogen, you have to take into consideration that there’s a lot more H-1 than H-2 and only a very tiny amount of H-3. That’s why the atomic mass of hydrogen on the periodic table isn’t a whole number: It’s 1.0079 amu. The number shows that there’s a lot more H-1 than H-2 and H-3.
Ions: Varying electrons
Because an atom itself is neutral, I say that the number of protons and electrons in atoms are equal throughout this book.
But in some cases, an atom can acquire an electrical charge. For example, in the compound sodium chloride — table salt — the sodium atom has a positive charge and the chlorine atom has a negative charge. Atoms (or groups of atoms) in which there are unequal numbers of protons and electrons are called ions.
The neutral sodium atom has 11 protons and 11 electrons, which means it has 11 positive charges and 11 negative charges. Overall, the sodium atom is neutral, and it’s represented like this: Na. But the sodium ion contains one more positive charge than negative charge, so it’s represented like this: Na+ (the + represents its net positive electrical charge).
Gaining and losing electrons
Atoms become ions by gaining or losing electrons. And ions that have a positive charge are called cations. The progression goes like this: The Na+ ion is formed from the loss of one electron. Because it lost an electron, it has more protons than electrons, or more positive charges than negative charges, which means it’s now called the Na+ cation. Likewise, the Mg2+ cation is formed when the neutral magnesium atom loses two electrons.
Now consider the chlorine atom in sodium chloride. The neutral chlorine atom has acquired a negative charge by gaining an electron. Because it has unequal numbers of protons and electrons, it’s now an ion, represented like this: Cl-. And because ions that have a negative charge are called anions, it’s now called the Cl- anion. (You can get the full scoop on ions, cations, and anions in Chapter 5, if you’re interested. This here’s just a teaser.)
Writing electron configurations
Here are some extra tidbits about ions for your chemistry reading pleasure:
✓ You can write electron configurations and energy level diagrams for ions. The neutral sodium atom (11 protons) has an electron configuration of 1s22s22p63s'. The sodium cation has lost an electron — the valence electron, which is farthest away from the nucleus (the 3s electron, in this case). The electron configuration of Na+ is 1s22s22p6.
✓ The electron configuration of the chloride ion (Cl-) is 1s22s22p63s23p6. This is the same electron configuration as the neutral argon atom. If two chemical species have the same electron configuration, they’re said to be isoelectronic. Figuring out chemistry requires learning a whole new language, eh?
✓ This section has been discussing monoatomic (one atom) ions. But polyatomic (many atom) ions do exist. The ammonium ion, NH4+, is a polyatomic ion, or specifically, a polyatomic cation. The nitrate ion, NO3-, is also a polyatomic ion, or specifically, a polyatomic anion.
Predicting types of bonds
Ions are commonly found in a class of compounds called salts, or ionic solids. Salts, when melted or dissolved in water, yield solutions that conduct electricity. A substance that conducts electricity when melted or dissolved in water is called an electrolyte. Table salt — sodium chloride — is a good example.
On the other hand, when table sugar (sucrose) is dissolved in water, it becomes a solution that doesn’t conduct electricity. So sucrose is a nonelectrolyte.
REMEMBER. Whether a substance is an electrolyte or a nonelectrolyte gives clues to the type of bonding in the compound. If the substance is an electrolyte, the compound is probably ionically bonded (see Chapter 5). If it’s a nonelectrolyte, it’s probably covalently bonded (see Chapter 6).