Compounds: Formulas, Nomenclature, Masses, and Models

Chemical Formulas Explained: How to Identify, Name, and Write Compounds

Why is CO₂ called carbon dioxide, but FeCl₃ is iron(III) chloride instead of iron trichloride? Why does CO₂ have the prefix di-, but CO becomes monoxide instead of monooxide?

If these naming rules seem confusing, you’re not alone.

Chemists follow specific conventions when naming compounds and writing chemical formulas. These rules help scientists clearly communicate what a substance contains and how its atoms are arranged.

Once you understand these rules, chemical formulas become much easier to read—and much less intimidating.

What You’ll Learn:

• How chemical formulas represent compounds
• How ionic compound names and formulas are determined
• How covalent compound names use numerical prefixes
• How molecular formulas and models show structure
• How to calculate molecular and formula mass

Key Takeaways

• Chemical formulas show composition. A chemical formula uses element symbols and subscripts to indicate the types and numbers of atoms in the smallest unit of a substance—either a molecule (covalent compounds) or a formula unit (ionic compounds).
• Ionic compounds form from charged ions. Metals and nonmetals combine as cations and anions, and their formulas are written so that the total positive and negative charges balance, usually using the simplest whole-number ratio.
• Ionic naming depends on ion type. Binary ionic compounds use cation name + anion ending in –ide, transition metals use Roman numerals to indicate charge, and polyatomic ions keep their established names.
• Covalent compounds use prefixes to show atom counts. Binary covalent compounds (nonmetal + nonmetal) are named with numerical prefixes (mono-, di-, tri-, etc.) that indicate the exact number of atoms in the molecule.
• Different formulas and models for molecules reveal different structural information. Molecular formulas show composition, while condensed formulas, structural formulas, and molecular models illustrate how atoms connect and arrange in three dimensions.

Chemical Formula

When you see a chemical formula, what information are you actually looking at?

A chemical formula uses element symbols and numerical subscripts to show the type and number of atoms present in the smallest unit of a substance. For covalent substances, that smallest unit is a molecule. For ionic substances, it is a formula unit—the simplest whole-number ratio of ions in the compound.

For example:

• H₂O tells us that one molecule contains two hydrogen atoms and one oxygen atom.
• CaCl₂ tells us that calcium chloride contains one calcium ion for every two chloride ions.

Notice that subscripts apply only to the symbol immediately before them, and the subscript “1” is never written. CaCl₂ means one calcium and two chlorines—not Ca₁Cl₂.

But formulas alone don’t tell the whole story. How do we know how atoms are connected? How do we name these compounds? And how do we determine their masses? To answer those questions, let’s look at the two major types of compounds separately: ionic and covalent.

Ionic Compounds

How can you recognize an ionic compound?

A strong clue is the presence of a metal as the first element in the formula, or a polyatomic cation such as ammonium (NH₄⁺). Ionic compounds consist of positively charged cations and negatively charged anions arranged in a continuous three-dimensional lattice. They do not form individual molecules.

Because compounds are electrically neutral, the total positive charge must equal the total negative charge. This idea governs everything about writing ionic formulas.

Naming Binary Ionic Compounds

How do we name the chemical formulas of ionic compounds? Let’s first look at the simplest kind of ionic compound: binary ionic compounds, which are composed of only monatomic ions.

To name binary ionic compounds:

1. Name the cation first.
2. Name the anion second.
3. Change the ending of the anion to -ide.

General rule for naming binary ionic compounds. The cation retains the name of the metal, while the ending of the nonmetal (anion) changes to “-ide.”

Hence, NaCl becomes sodium chloride and CaO becomes calcium oxide.

Formulas of Ionic Compounds

How do you determine the chemical formula just from the name of the compound?

Suppose you are given calcium bromide. First, identify the ions:

• Calcium (Group 2) forms Ca²⁺.
• Bromine (Group 17) forms Br^–.

Next, balance the charges. One Ca²⁺ requires two Br^– ions to neutralize the +2 charge. Hence, the formula becomes: CaBr₂. This charge-balancing principle explains why aluminum oxide is Al₂O₃. Aluminum forms Al³⁺ and oxygen forms O^2–. Two Al³⁺ ions have a total charge of +6, so it balances out with three O^2– ions, which have a total charge of −6.

Alternatively, use the crisscross or crossover method—the charge of one ion becomes the subscript of the other element (the signs of the charges are dropped in the chemical formula).

Crossover method. A shortcut to write the formula of ionic compounds without calculating the charge balance is to write the charge of the other ion as the subscript of the other.

The chemical formulas for ionic compounds are always written with the ion ratios in their simplest whole-number form. In other words, when charges are “crossed over” to determine subscripts, the resulting ratio should be reduced if possible.

Simplifying ion ratios. Ionic formulas are written using the lowest whole-number ratio of cations and anions needed to achieve charge balance.

For example, calcium oxide is formed from Ca²⁺ and O^2– ions. Crossing over the charges would initially give Ca₂O₂, but this ratio simplifies to CaO, which is the correct formula.

This rule may appear to have a few exceptions, particularly for compounds containing the Hg₂²⁺ ion, such as Hg₂Cl₂, and for peroxides, which contain the O₂^2– ion, such as Na₂O₂.

However, these are not true exceptions. In these cases, the Hg₂²⁺ and O₂^2– ions already exist as paired atoms that function as a single polyatomic ion. Because the atoms are already bound together within the ion, they cannot be separated when forming compounds.

As a result, when these ions combine with other ions—especially those whose charge magnitude is an odd number, such as +1 or −1—the subscripts represent the number of entire ions, not individual atoms. Simplifying the subscripts further would incorrectly break apart the polyatomic ion.

Ionic Compounds with Transition Metal Ions

What happens if a metal can form more than one charge? This is where transition metals require special attention.

Many transition metals can lose different numbers of electrons, forming cations with different charges. Iron, for example, commonly forms Fe²⁺ and Fe³⁺. Copper forms Cu⁺ and Cu²⁺. Because these ions produce different compounds with different properties, their names must clearly distinguish between them.

There are two systems used to name transition metal cations.

Table 1. Common Transition Metal Ions That Can Have Multiple Charges

Metal	Ion Formula	IUPAC Name	Common Name
Copper	Cu+	Copper(I)	Cuprous
	Cu2+	Copper(II)	Cupric
Gold	Au+	Gold(I)	Aurous
	Ag3+	Gold(III)	Auric
Iron	Fe2+	Iron(II)	Ferrous
	Fe3+	Iron(III)	Ferric
Chromium	Cr2+	Chromium(II)	Chromous
	Cr3+	Chromic(III)	Chromic
Cobalt	Co2+	Cobalt(II)	Cobaltous
	Co3+	Cobalt(III)	Cobaltic
Lead	Pb2+	Lead(II)	Plumbous
	Pb4+	Lead(IV)	Plumbic
Tin	Sn2+	Tin(II)	Stannous
	Sn4+	Tin(IV)	Stannic
Mercury	Hg22+	Mercury(I)	Mercurous
	Hg2+	Mercury(II)	Mercuric

The Modern IUPAC System

The simpler and more modern method is the IUPAC system, developed by the International Union of Pure and Applied Chemistry. In this system, the charge of the metal ion is written as a Roman numeral in parentheses immediately after the metal’s name—there is no space between the name of the metal and the Roman numeral in parentheses.

For example:

• Fe²⁺ → iron(II) ion
• Fe³⁺ → iron(III) ion
• Cu⁺ → copper(I) ion
• Cu²⁺ → copper(II)

Let’s use the crossover method to write the formula for a binary ionic compound that contains a transition metal cation, iron(II) chloride. Recall that the Roman numeral pertains to the charge of the metal ion, not the number of metal atoms in the formula unit. Hence, iron(II) chloride consists of:

• Fe²⁺ ion (not two Fe atoms)
• Cl^– ion

Crossing over the charges (and dropping the signs) gives FeCl₂.

Reversing the crossover method. By applying the crossover method and the principle of charge balance, the charge of a cation with multiple possible charges can be determined.

Likewise, the crossover method can be used in reverse to determine the name of an ionic compound with a transition metal. For instance, since Br has no subscript in CuBr, the copper must have a charge of +1, and the compound would be named copper(I) bromide.

The Common (Classical) System

Before the IUPAC system was established, chemists used what is called the common system or classical system. This system uses two suffixes:

• -ous → lower relative charge
• -ic → higher relative charge

For example:

• Co²⁺ → cobaltous ion
• Co³⁺ → cobaltic ion

For metals that have their Latin names on the periodic table, the Latin name is often used. For example, iron is ferrum, so:

• Fe²⁺ → ferrous ion
• Fe³⁺ → ferric ion

Likewise, copper is cuprum, so:

• Cu⁺ → cuprous ion
• Cu²⁺ → cupric ion

IUPAC vs. common system naming of ionic compounds. For cations that can have multiple charges, either the IUPAC or the common system must be used although the IUPAC system is now preferred.

The use of the IUPAC system is recommended and now the standard, but the classical system names are still chemically correct and commonly encountered in medicine and the health sciences.

Why the IUPAC System is Preferred

Although widely used historically, the common system has two major limitations.

1. It can only distinguish two charges. Because it uses only –ous and –ic, it works only for metals that form exactly two ions. It cannot properly distinguish metals like vanadium that form four different ions:

• V²⁺ → vanadium(II) ion
• V³⁺ → vanadium(III) ion
• V⁴⁺ → vanadium(IV) ion
• V⁵⁺ → vanadium(V) ion

2. It shows relative charge, not exact charge. The suffix only tells you which charge is higher or lower—not what the actual charge is. For example:

• Iron’s lower charge is +2 → ferrous
• Copper’s lower charge is +1 → cuprous

Both use -ous for the lower charge, but the numerical charges are different. This inconsistency frequently confuses students. Even more confusing:

• Fe²⁺ ends in -ous
• Cu²⁺ ends in -ic

So the same +2 charge does not always use the same suffix across different metals.

For these reasons, the IUPAC Roman numeral system is strongly preferred in modern chemistry—it removes ambiguity and directly communicates the ion’s charge.

Note that these naming systems only apply to metal ions that can have multiple charges, so they are not used in main-group elements—Na⁺, Ca²⁺, and Al³⁺ are still sodium, calcium, and aluminum ions respectively. They are not named sodium(I), calcium(II), and aluminum(III).

Polyatomic Ions

Recall that polyatomic ions are charged groups of atoms that stay together and behave as a single unit in chemical reactions and formulas. Hence, when writing formulas, polyatomic ions must remain intact. When naming compounds, the polyatomic ion keeps its established name.

Two common examples of polyatomic ions are nitrate (NO₃^–) and sulfate (SO₄^2–). If Fe³⁺ combines with nitrate, charges are balanced as before: Fe³⁺ requires three NO₃^– ions. The formula becomes: Fe(NO₃)₃.

Why are parentheses used? Because there are two entire nitrate ions. If only one were present, as in NaNO₃, parentheses would not be needed.

Examples of ionic compounds with polyatomic ions. Polyatomic ions retain their names when naming compounds.

Another common polyatomic ion is the cation ammonium, NH₄⁺. It is essential not to confuse ammonium (NH₄⁺), which has a positive charge, with ammonia (NH₃), which is a neutral molecule and contains one less hydrogen atom. That one extra hydrogen in ammonium makes all the difference in charge.

Table 2. Some common polyatomic ions

Common polyatomic ions

Most polyatomic ions you will study, however, are anions, and many of them contain oxygen. These oxygen-containing polyatomic anions are called oxoanions.

Although many polyatomic ion names must be memorized, oxoanions have clear naming patterns that make them much easier to learn.

Oxoanion Pairs

When an element forms two oxoanions that differ only in the number of oxygen atoms, the naming pattern is straightforward:

• The ion with more oxygen atoms ends in -ate.
• The ion with fewer oxygen atoms ends in -ite.

For example:

• NO₃^– → nitrate
• NO₂^– → nitrite
• SO₄^2– → sulfate
• SO₃^2– → sulfite

If you remember that “ate” means more oxygen than “ite,” you can often predict the correct name.

Larger Oxoanion Families

Some elements, especially the halogens, form four related oxoanions. How do we distinguish them?

In these cases:

• The ion with the most oxygen atoms uses per- … -ate
• The next uses -ate
• The next uses -ite
• The ion with the fewest oxygen atoms uses hypo- … -ite

Naming for families with four oxoanions. Only the number of oxygen atoms varies while the charge remains the same. In addition to the -ate and -ite suffixes, prefixes (per- and hypo-) are used for the ions with the most and fewest oxygen atoms.

In other words, as oxygen content increases, the naming progresses from hypo-ite → ite → ate → per-ate. Chlorine provides the classic example:

• ClO₄^– → perchlorate
• ClO₃^– → chlorate
• ClO₂^– → chlorite
• ClO^– → hypochlorite

This systematic pattern prevents random memorization and replaces it with logic.

Hydrogen-Containing Oxoanions

What happens if hydrogen is part of the polyatomic anion itself? This occurs when a hydrogen ion (H⁺) is added to make a new polyatomic ion. In this case, the word “hydrogen” is added before the oxoanion name. For example:

• HCO₃^– → hydrogen carbonate
• HSO₄^– → hydrogen sulfate

Adding a H⁺ reduces the magnitude of negative charge. Carbonate and sulfate have the formulas CO₃^2– and SO₄^2–, respectively. The added H⁺ partially neutralizes their negative charge, so the –2 charge becomes –1 in HCO₃^– and HSO₄^–.

If two hydrogens are present, the prefix “dihydrogen” is used:

• H₂PO₄^– → dihydrogen phosphate

Phosphate originally has the formula PO₄^3–. Since two H⁺ ions are added, the charge changes by +2, converting the –3 charge of PO₄^3– into the –1 charge of H₂PO₄^–.

From these examples, a pattern emerges: each hydrogen added to an oxyanion increases the charge by +1 (or reduces the negative charge by one unit).

Effect of hydrogen on the charge of oxoanion. For each hydrogen atom added, the charge of an oxoanion increases by one.

Some hydrogen-containing oxoanions also have common names that are still widely used. For example, HCO₃^– (hydrogen carbonate) is commonly called bicarbonate.

Understanding this pattern is important because these ions appear frequently in biological systems, environmental chemistry, and laboratory reactions. These patterns also make memorization much easier.

Summary for Deriving Ionic Formulas and Names

Let’s sum up how ionic compounds are named, and how formulas can be derived from these names.

Formula → Name: When you are given a formula, what should you do first?

1. Identify ion names and their charges.
2. Name the cation and indicate charge if needed.
   1. If it is a main-group metal, use the element name.
   2. If it is a transition metal, use Roman numerals to indicate the charge.
   3. If it is a polyatomic cation, use the ion name.
3. Name the anion.
   1. If it is monatomic, change ending to -ide
   2. If it is polyatomic, use ion name

Name → Formula: When given a name, what must you do?

1. Identify ion symbols and charges.
2. Balance charges (use crossover method).
3. Use smallest whole-number ratio.
4. Add parentheses if needed.

Hydrates: Compounds with Water

Some ionic compounds include a specific number of water molecules—called waters of hydration—trapped in their structure. These compounds are called hydrates. Their formulas include a centered dot, and they are named by the name of the ion plus a Greek numerical prefix indicates the number of water molecules.

Table 3. Numerical prefixes for hydrates and binary covalent compounds

Number	Prefix
1	mono-
2	di-
3	tri-
4	tetra-
5	penta-
6	hexa-
7	hepta-
8	octa-
9	nona-
10	deca-

For example:

• MgSO₄·7H₂O → magnesium sulfate heptahydrate (commonly known as Epsom salt)
• Na₂CO₃·10H₂O → sodium carbonate decahydrate (commonly known as washing soda)

Heating can remove these waters of hydration, producing a different compound.

Acids: Special Hydrogen Compounds

Acids are one of the most important classes of hydrogen-containing compounds in chemistry. When naming acids and writing their formulas, we treat acids as anions with one or more hydrogen ions (H⁺) added to make a neutral compound. As such, their naming naturally follows ionic compound conventions.

There are two main types of acids: binary acids and oxoacids or oxyacids.

Binary Acids

What makes an acid “binary”? As with all binary compounds, binary acids contain only two distinct elements: hydrogen and one other nonmetal. What distinguishes them from oxoacids is that binary acids do not contain oxygen.

Binary acids often begin as gaseous hydrogen compounds that form acidic solutions when dissolved in water. Their names follow this pattern:

hydro– + nonmetal root + –ic + acid

For example, hydrogen chloride gas (HCl) is a covalent compound. However, when it dissolves in water, it ionizes to form H⁺ and Cl^– ions, producing an acidic solution. In aqueous solution, it is therefore named hydrochloric acid.

Naming binary acids. Binary acids form when gaseous covalent compounds form ions in water. They contain hydrogen, another nonmetal, and no oxygen.

This naming pattern reflects the change in behavior: the compound exists as a molecular substance in the gas phase, but forms ions in water, which gives the solution its acidic properties.

This pattern works for hydrogen combined with anions that normally end in -ide:

• Cl^– → chloride → hydrochloric acid
• Br^– → bromide → hydrobromic acid
• S^2– → sulfide → hydrosulfuric acid

Notice the pattern:

• Anion ends in -ide → acid ends in -ic acid
• Prefix hydro- is added

Oxoacids (Oxyacids)

What makes an acid an oxoacid? An oxoacid contains oxygen in addition to hydrogen and another element as they are derived from oxoanions. Consequently, the name of an oxoacid is directly based on the name of its oxoanion.

How do oxoacid names relate to oxoanions? Two simple suffix changes occur:

• -ate → -ic acid (oxoanion has more O atoms)
• -ite → -ous acid (oxoanion has less O atoms)

Naming oxoacids. Oxoacids form when oxoanions combine with hydrogen ions.

What about “per-” and “hypo-” oxoanions? The prefixes used in oxoanions are retained in oxoacid names. For example:

• BrO₄^–, perbromate → HBrO₄, perbromic acid
• ClO^–, hypochlorite → HClO, hypochlorous acid

Covalent Compounds

Covalent compounds consist of discrete molecules formed when nonmetals bond together by sharing electrons. Because these substances exist as individual molecules rather than extended lattices, their chemical formulas are specifically called molecular formulas.

Many covalent compounds are also organic compounds. Because organic chemistry includes many different subclasses of molecules with their own specialized naming rules, their nomenclature is typically discussed in a separate study guide. In this section, we focus on the naming and formula writing of binary covalent compounds, which contain two different nonmetal elements.

Naming Binary Covalent Compounds

One key difference between naming ionic and covalent compounds is that the names of molecular formulas show the actual number of atoms in each molecule using numerical prefixes (Table 3).

• The first element uses a prefix only if more than one atom is present
• The second element almost always uses a prefix and ends in -ide.

For example, N₂F₄ is named dinitrogen tetrafluoride and NF₃ is nitrogen trifluoride. Since only one N atom is present in NF₃ and it is the first element, the prefix “mono-” is not used. However, if there is only one atom for the second element, the prefix is used. CO, for instance, is carbon monoxide; not monocarbon monoxide.

Rules for naming binary covalent compounds. Numerical prefixes indicate the number of atoms of each element. The first element retains its name, while the second ends in -ide.

The name for CO also shows another naming convention: when prefixes ending in a vowel combine with “oxide,” a vowel is often dropped for easier pronunciation. For example monooxide becomes monoxide, and tetraoxide becomes tetroxide.

These naming rules imply the significance of the order of elements in the formula, but how do we know which element to write and name first?

When naming and writing molecular formulas, the order of elements in the name generally follows their positions in the periodic table:

• The element with the lower group number is named first. For example, in compounds formed between nitrogen (Group 15) and fluorine (Group 17), nitrogen is always written first: NF₃ (nitrogen trifluoride) and N₂F₄ (dinitrogen tetrafluoride).

• If the two elements belong to the same group, the element in the higher period (lower in the periodic table) is named first. For instance, sulfur (Period 3) and oxygen (Period 2) form SO₂, which is named sulfur dioxide.

There are notable exceptions to this ordering rule. When oxygen combines with the halogens chlorine, bromine, or iodine, the halogen is named first. Examples include dichlorine monoxide (Cl₂O) and diiodine pentoxide (I₂O₅). Another exception arises from tradition—notice that the formula for ammonia is written as NH₃ instead of H₃N.

Ammonia also illustrates another effect of tradition: some covalent compounds are so common that they are known by traditional or common names rather than systematic ones. Other examples include H₂O, commonly called water instead of dihydrogen monoxide, and HCl, called hydrogen chloride instead of hydrogen monochloride. However, most binary covalent compounds are named systematically using the rules discussed above.

Formulas of Binary Covalent Compounds

Writing the formula for a binary covalent compound involves reversing the naming process. Because covalent compounds do not form ions, there is no charge balancing step when writing formulas for covalent compounds, and the prefixes in the name directly determine the number of atoms in the formula. As such, simply ask: what do the prefixes tell us about the number of atoms?

Each prefix corresponds to a specific number, which becomes the subscript in the chemical formula.

For example:

• Sulfur dioxide

Sulfur → S

Di-oxide → O₂

Formula: SO₂

• Dinitrogen trioxide

Di-nitrogen → N₂

Tri-oxide → O₃

Formula: N₂O₃

• Carbon tetrachloride

Carbon → C

Tetra-chloride → Cl₄

Formula: CCl₄

• Diphosphorus pentoxide

Di-phosphorus → P₂

Penta-oxide → O₅

Formula: P₂O₅

Unlike in the formulas for ionic compounds, the subscripts in molecular formulas are not written with their smallest whole-number ratios since covalent formulas reflect the actual number of atoms in a molecule. For instance, the molecular formula of glucose is always written as C₆H₁₂O₆, not CH₂O.

Summary for Deriving Molecular Formulas and Names

Let’s now summarize how binary covalent molecular compounds are named and how their formulas can be derived from those names.

Formula → Name: When you are given a formula, what should you do first?

1. Identify the elements present.
2. Name the first element.
   1. Use the element name.
   2. Use a prefix only if more than one atom is present.
3. Name the second element.
   1. Use a prefix to indicate the number of atoms.
   2. Change the ending of the element name to -ide.
   3. Adjust for “oxide” if necessary—when prefixes ending in vowels combine with “oxide,” a vowel may be dropped (for example, monoxide instead of monooxide).

Name → Formula: When you are given a name, what must you do?

1. Identify the elements mentioned in the name.
2. Write their chemical symbols in the order given.
3. Convert prefixes into subscripts. Do not reduce subscripts to their smallest ratios.

Naming for ionic and covalent compounds. This flowchart summarizes how the name of a compound can be determined from its chemical formula. 

Molecular Formulas and Models

In addition to molecular formulas, covalent compounds can be represented using several types of formulas and models that convey different kinds of structural information.

A molecular formula shows the types and numbers of atoms present in a molecule, but it does not show how those atoms are connected or arranged in space.

This limitation becomes especially important in organic chemistry, where different compounds can share the same molecular formula but have different structures. For example, ethanol and dimethyl ether both have the molecular formula C₂H₆O, yet their atoms are connected differently. Because of these different structures, the two compounds have different properties.

Formulas and models to represent molecules. Formulas show how atoms are connected and the number of each atom present, while models illustrate the three-dimensional arrangement of atoms.

To highlight how atoms are connected within a molecule, formulas are used:

• Condensed formulas. These are commonly used for organic compounds. They show each carbon atom and the atoms attached to it, but they do not explicitly draw every bond.
• Structural formulas. These formulas explicitly show how atoms are bonded by connecting them with lines that represent covalent bonds.

On the other hand, molecular models help visualize the three-dimensional shape of molecules.

• Ball-and-stick models. Atoms are represented as balls and bonds as sticks. These models clearly show bond angles and molecular geometry, but they often exaggerate the distances between atoms.
• Space-filling models. These models represent atoms as spheres whose sizes correspond more closely to their relative atomic radii. Space-filling models show how closely atoms pack together in a molecule, although individual atoms can become harder to distinguish in complex molecules.

Each representation provides different information, and no single model tells the whole story.

Let’s help you remember the key topics with these memory tricks or mnemonics

“C-A-IDE”

Naming monatomic ions in ionic compounds:

• Cation first (metal name unchanged)
• Anion second (-ide ending)

“ic = high, ous = low”

Common system naming for ions with multiple charges:

• Ionic compounds that contain the cation with the higher charge use –ic as the metal suffix.
• If it contains the cation with the lower charge, the suffix is –ous.

Oxoacids

• If the acid is composed of the oxoanion with more oxygen, the –ate suffix becomes –ic.
• If the acid is composed of the oxoanion with less oxygen, the –ite suffix becomes –ous.

“Ate more oxygen, ite had less.”

Names of oxoanions

• The ion with the more oxygen ends in –ate.
• The ion with the less oxygen ends in –ite.

“1 H, +1 charge.”

Hydrogen in oxoanions

• The charge of an oxoanion increases by +1 with every H atom added.

“Lower group goes first; lower row is the alternative.”

Order of writing atoms in molecular formula

• The element with lower group number is written and named first in covalent compounds.
• If both elements belong to the same group, the element from the lower row comes first.

Conclusion

Now the questions from the beginning should make much more sense.

Why is CO₂ called carbon dioxide, but FeCl₃ is iron(III) chloride instead of iron trichloride? The answer lies in the type of compound involved. Covalent compounds use numerical prefixes to show the number of atoms in a molecule, while ionic compounds follow charge-balancing rules and may require Roman numerals to indicate the charge of a metal ion.

Throughout this guide, you’ve seen how chemists use a consistent system to describe substances. Chemical formulas reveal which atoms are present and how many there are. Naming rules allow scientists to translate formulas into words and names back into formulas, and molecular formulas and models help visualize how atoms connect and arrange in space.

Together, these tools form the language of chemistry—a clear and systematic way to describe how matter is built from atoms.

Once you recognize the patterns behind formulas, prefixes, and naming conventions, chemical formulas stop looking like random symbols. Instead, they become useful descriptions that tell you exactly what a compound contains and how it is structured.