When you first start learning how to name ionic or molecular compounds, a common strategy is to try and memorize everything. You can do this, but it’s a lot of work, and there is an easier approach. Instead, understand that the concept behind naming is that all ionic compounds form with cation to anion ratios that give neutral molecules. These mole ratios are indicated in the name, as well as shown as subscripts in the formula. While molecular (covalent) compounds do not form to balance out charges like ionic compounds do, the ratio of elements present is still reflected in the name and formula.
Organization of the Periodic Table
Before you dive into naming, it’s important to have a thorough understanding of periodic trends, and how/why different elements behave the way they do. The periodic table is organized into columns known as groups or families based on their shared number of valence electrons. Elements in the same group have similar chemical behaviors, meaning they gain or lose the same number of electrons to achieve an octet in their outermost shell. The groups are labeled from 1 to 18.
Elements in Group 1, the alkali metals, possess a single electron in their outermost shell. Loosing 1 valence electron to achieve an inner shell octet is much easier than gaining 7 electrons to fill and outer shell. This means that elements in Group 1 become cations with a charge of +1 when they ionize. Group 2 elements, alkali earth metals, have two outer electrons and become +2 cations when they ionize for the same reason. Moving across the table, the transition metals occupy the central region (groups 3-12). Transition metals are characterized by their variable oxidation states, meaning their charges can vary depending on what type of compound they form. Elements in Group 13 have 3 valence electrons which they give up when they ionize, becoming a +3 cation. Notice a trend? Every time you move one tall column to the right, the number of valence electrons increases by one. This means that Group 14 has 4 valence electrons. Once you hit Group 15, the way elements ionize changes. These elements have 5 valence electrons. It is now easier to gain 3 electrons and fill their outer shell rather than to loose 5 to achieve an octet. Elements in Group 15 become anions with a charge of -3. Following the left to right trend, Group 16 has 6 valence electrons. Elements in this group gain 2 electrons to fill their outermost shells and become anions with a -2 charge. Group 17 contains the halogens which are highly reactive nonmetals. They gain one electron to become -1 anions. The noble gases are in Group 18. These elements are rarely included in forming compounds because they are stable on their own due to their complete outer electron shells.
Metals vs. Non-Metals
Now that you have a refresher on the general organization of the table, you should have a clear idea of the difference between metals and non-metals. Metals are situated predominantly on the left and center of the periodic table, and share the common characteristic of losing electrons to form positively charged ions (cations). In contrast, non-metals are located to the right of the stair-step of the periodic table. They gain electrons and become negative anions to achieve a stable electron configuration.
Ionic compounds are formed through the transfer of electrons. When a metal and a non-metal come into contact, the metal atom donates one or more electrons to the non-metal atom. This electron transfer leads to the formation of ions with opposite charges—positively charged cations from the metal and negatively charged anions from the non-metal. The attractive electrostatic forces between these oppositely charged ions result in the formation of an ionic bond, holding the ions together in a crystal lattice structure.
The Basics of Naming Ionic Compounds
To start, we will name ionic compounds only using elements from the tall columns on the periodic table (Groups 1,2 and 13-18), which have defined charges. Anytime the name of an ionic compound is given, you should be able to determine the ratio of ions in that compound based on balancing their charges and also write a correct molecular formula. There are lots of different rules for naming that apply to specific circumstances, but the basic rules for naming used for tall column elements apply to everything. First, when naming, always name the metal first. For example, when naming a compound containing lithium (a metal) and chlorine (a non-metal), you would say “lithium chloride” with the lithium being named first. Even in situations where there is no metal present (covalent compounds), the most metallic is named first. “Most metallic” means the non-metal in the compound that is closest to the metals on the periodic table. Second, the non-metal gets an “-ide” ending. This ending tells you that the non-metal ion is monatomic (vs. polyatomic).
Naming Ionic Compounds Containing Transition Metals
If you’ve heard of a metal before you took chemistry, it’s probably a transition metal. For example, silver, gold, aluminum, iron, tin, etc.. We said before that the charges on groups 3-12 (transition metals in the short columns) can vary. They lose different numbers of electrons to form positive ions with different charges. If this is the case, the specific charge the metal takes on must be indicated in the name. Roman numerals in parentheses are used to indicate the oxidation state of the transition metal in the compound. For example, iron (Fe) can form ions with a +2 or +3 charge by losing either 2 or 3 electrons respectively. For instance, iron(II) chloride (FeCl2) contains iron with a +2 charge, and iron(III) chloride (FeCl3) contains iron with a +3 charge. Again, the Roman numeral tells the charge on the transition metal, NOT the number of metals present in the compound. The transition metal with its Roman numeral is named first, and monatomic non-metals in compounds containing transition metals are given the “-ide” ending.
Naming Ionic Compounds Containing Polyatomic Ions
A polyatomic ion is a charged species composed of two or more non-metal atoms covalently bonded together. Unlike simple ions (monatomic ions), which consist of a single atom with a positive or negative charge, polyatomic ions are groups of elements that carry an overall charge. Common examples include the sulfate ion (SO₄²⁻), which consists of one sulfur atom and four oxygen atoms, and the ammonium ion (NH₄⁺), composed of one nitrogen atom and four hydrogen atoms. Sometimes teachers will allow you to use a given list of polyatomic ions on tests. Even so, it is important to memorize the ones that your class uses frequently. If you just rely on the given list, chances are that you will not recognize polyatomic ions in compounds. This leads to a lot of problems. Polyatomic ions can have either a positive or negative charge, and their chemical behavior in ionic bonding is the same as that of monatomic ions with the same charge. In essence, you treat the whole group the same as you would treat a monatomic ion when balancing charges in an ionic compound. For example, because the sulfate ion mentioned above has a -2 charge, it will behave the same as the Group 16 non-metals when forming compounds. That said, they are not named in the same way. If a polyatomic ion is present in a compound, the metal is named first, and then the polyatomic ion is named using the “-ate” or “-ite” ending. In cases where a compound contains both a positive and negative polyatomic ion, the cation is written first, followed by the name of the polyatomic anion.
Naming Covalently Bonded Molecular Compounds
So far, all of the naming rules have applied to scenarios based on charge balancing. Molecular compounds form differently. Rather than electrons being transferred and forming balanced cation/anion units, in covalent compounds, electrons are shared between non-metals. The names of these compounds are based on the systematic use of prefixes to indicate the number of atoms of each element in the molecular formula. Prefixes include "mono-" for one, "di-" for two, "tri-" for three, "tetra-" for four, and so on. However, the prefix "mono-" is omitted for the first element. For example, CO2 is named carbon dioxide, not monocarbon dioxide. The element that appears first in the formula is the one with the lower electronegativity (the most metallic), and its name is written first. The second element's name is then modified to end in "-ide."
It's important to note that some compounds have common names that are widely accepted in addition to their systematic names. For example, H2O is commonly known as water, and NH3 is ammonia. These common names are often used in everyday language and are universally recognized.