# 4.6: Ionic Formula Writing

• • Contributed by Elizabeth Gordon
• Lecturer (Chemistry) at Furman University

Learning Objectives

• Write formulas for binary ionic compounds
• Relate the Lewis structure to the chemical formula for binary ionic compounds.
• Commit bolded polyatomic formulas to memory.
• Write formulas for ternary (polyatomic) ionic compounds.

## Ionic Formulas

Chemical formulas for ionic compounds are called ionic formulas. A proper ionic formula has a cation and an anion in it; an ionic compound is never formed between two cations only or two anions only. The key to writing proper ionic formulas is simple: the total positive charge must balance the total negative charge. Because the charges on the ions are characteristic, sometimes we have to have more than one of a cation or an anion to balance the overall positive and negative charges. It is conventional to use the lowest ratio of ions that are needed to balance the charges. Figure $$\PageIndex{1}$$: Image taken from: https://upload.wikimedia.org/wikiped...my_Bonding.JPG

For example, consider the ionic compound between $$\ce{Na^{+}}$$ and $$\ce{Cl^{−}}$$. Each ion has a single charge, one positive and one negative, so we need only one ion of each to balance the overall charge. When writing the ionic formula, we follow two additional conventions:

1. write the formula for the cation first and the formula for the anion next, but
2. cross the charges down diagonally.

Thus, for the compound between $$\ce{Na^{+}}$$ and $$\ce{Cl^{−}}$$, we have the ionic formula $$\ce{NaCl}$$ (Figure $$\PageIndex{1}$$). The formula $$\ce{Na2Cl2}$$ also has balanced charges, but the convention is to use the lowest ratio of ions, which would be one of each. (Remember from our conventions for writing formulas that we do not write a 1 subscript if there is only one atom of a particular element present.) For the ionic compound between magnesium cations ($$\ce{Mg^{2+}}$$) and oxide anions ($$\ce{O^{2−}}$$, again we need only one of each ion to balance the charges. By convention, the formula is $$\ce{MgO}$$. Figure $$\PageIndex{2}$$: NaCl is the formular for Table Salt. The ionic compound NaCl is very common. Red rock salt from the Khewra Salt Mine in Pakistan. Image used with permission (CC BY-SA 4.0; Hubertl).

Rules for Writing Ionic Formulas

Recall ionic compound can be of the following combination: metal/nonmetal or ammonium/nonmetal

1. Recall the cation charge. Superscript this charge to the right of the cation (Ca2+). If metal is multicharged metal (like most transitions and metals below metalloid staircase), then charge is indicated by roman numeral (Lead IV would translate to Pb4+).
2. Recall the anion charge. Superscript this charge to the right of the anion (S2-). If polyatomic ion is present, then put poly in parentheses and superscript charge to the right of this species.
3. Cross charges down diagonally and lose signs. Subscripts that are divisible by the same number must be reduced (ionics are simple compounds). Do not change what is inside a poly. You may keep your parentheses around polyatomic ions.

* Binary compound names will end with ide (except cyanide and hydroxide). This means the anion will come from the periodic table.

** Compounds that contain polyatomic ions will end in ate/ite or the two weird ides (cyanide/hydroxide). You will need to memorize the polyatomic ion table below.**

Example $$\PageIndex{1}$$:

Write the formulas for the combinations of ions below:

1. Ca2+ and Cl
2. Al3+ and F
3. Li+ and O2−

Solution

1. The proper ionic formula is CaCl2.
2. The formula for this compound is AlF3.
3. After crossing charges diagonally down, the formula will be Li2O.

Exercise $$\PageIndex{2}$$

Write the proper ionic formulas for each of the two given ions.

1. Fe2+ and S2−
2. Fe3+ and S2−

$$\ce{FeS}$$, Subscripts must be reduced for ionic compounds. If this is not done, then the answer will be incorrect.

$$\ce{Fe2S3}$$, These subscripts cannot be reduced because they are not divisible by the same number.

There also exists a group of ions that contain more than one atom. These are called polyatomic ions. Table $$\PageIndex{1}$$ lists the formulas, charges, and names of some common polyatomic ions. Only one of them, the ammonium ion, is a cation; the rest are anions. Most of them also contain oxygen atoms, so sometimes they are referred to as oxyanions. Some of them, such as nitrate and nitrite, and sulfate and sulfite, have very similar formulas and names, so care must be taken to get the formulas and names correct. Note that the -ite polyatomic ion has one less oxygen atom in its formula than the -ate ion but with the same ionic charge.

Table $$\PageIndex{1}$$: Common Polyatomic Ions
Name Formula and Charge Name Formula and Charge
ammonium (NH4)+ nitrate (NO3)-
acetate (C2H3O2)- nitrite (NO2)-
bicarbonate (hydrogen carbonate) (HCO3)- carbonate (CO3)2-
hydroxide (OH)- sulfate (SO4)2-
chlorate (ClO3)- sulfite (SO3)2-
hypochlorite (ClO)- phosphate (PO4)3-
cyanide (CN) phosphite (PO3)3-

You will need to commit Table $$\PageIndex{1}$$ to memory for future tests and quizzes. There are more polyatomic ions than the ones listed in the table above. You will be provided with a list of these additional polyatomic ions for use on tests/quizzes. In our class, this list (List O' Polyatomics) will have the ones you are required to memorize in bold print.

Writing the formulas of ionic compounds has one important difference. If more than one polyatomic ion is needed to balance the overall charge in the formula, enclose the formula of the polyatomic ion in parentheses and write the proper numerical subscript to the right and outside the parentheses. Thus, the formula between calcium ions, Ca2+, and nitrate ions, NO3, is properly written Ca(NO3)2, not CaNO32 or CaN2O6. In our course, I encourage you to keep your polyatomics insides parentheses at all times. Refer back to your rules for writing ionic formulas.

Example $$\PageIndex{4}$$:

After writing these formulas, watch your instructor teach more examples of this concept by clicking on this link.

1. ammonium sulfide
2. aluminum phosphate
3. iron II phosphite

Solution

1. Ammonium is a polyatomic ion (cation). These species will go first in the formula. Take ammonium (NH4)+ and combine it with sulfide S2-(ide = periodic table, except for hydroxide/cyanide). While keeping the polyatomic formula within parentheses, cross the ion charges down and lose signs. The resulting formula will be (NH4)2S.
2. Aluminum is Al3+ and phosphate is (PO4)3-. While keeping the polyatomic ion inside parentheses, cross charges down and reduce subscripts outside. The final formula for the compound will be Al(PO4). For more advanced chemistry courses, students are encouraged to drop the parentheses if a subscript is not present once charges have been crossed down. I am fine with you always keeping parentheses.
3. Transition metal charges (except for silver and zinc) will always be noted by a Roman numeral. This charge is always positive for the metal ion. Place this charge to the top right corner of the metal atom, Fe2+. Next, recall that phosphite is the polyatomic, (PO3)3-. Cross the charges down diagonally and keep the polyatomic inside the parentheses. The final result will be Fe3(PO3)2.

Exercise $$\PageIndex{4}$$

Write the formulas for the following ionic compounds:

1. ammonium bicarbonate
2. cobalt III nitrite
4. potassium hydroxide
5. silver phosphide
6. mercury I sulfite

$$\ce{(NH4)(HCO3)}$$, there are no numbers outside the parentheses so you could write $$\ce{NH4HCO3}$$ for the answer. Do not combine the hydrogen atoms in the formula.

$$\ce{Co(NO2)3}$$, the parentheses must be kept for this answer.

$$\ce{PbO2}$$, after the charges have been crossed, subscripts must be reduced by the division of common factor

$$\ce{K(OH)}$$, you could drop ()

$$\ce{Ag3P}$$, phosphide is from the periodic table, not the polyatomic table. Know your endings!!

$$\ce{Hg2(SO3)}$$, you could drop ()