Lesson 4

Chemical Nomenclature

Learn systematic naming of ionic compounds, molecular compounds, acids, and polyatomic ions using IUPAC rules.

10 learning objectivescompounds naming
Chemistry reference tables

Ionic vs. Molecular Compounds

Many introductory compounds can be classified by how their atoms are bonded:

  • Ionic compounds consist of positively and negatively charged ions held together by electrostatic attraction. They typically have high melting points and conduct electricity when molten or when they dissolve to produce mobile ions. Examples: NaCl, CaO, MgBr⁠2.
  • Molecular (covalent) compounds consist of discrete molecules whose atoms share electrons. Many have lower melting points than ionic solids, and solutions of nonelectrolyte molecules conduct poorly. Examples: H⁠2O, CO⁠2, CH⁠4.
  • Network-covalent solids contain covalent bonds extending throughout the solid rather than discrete molecules. SiO⁠2 is a central example.

A useful first check is composition: metal cations paired with nonmetals or polyatomic anions usually indicate an ionic compound, while discrete compounds containing only nonmetals are usually molecular. This is a heuristic, not a law. NH⁠4Cl is ionic despite containing no metal, and anhydrous AlBr⁠3 is covalent despite containing aluminum.

Go deeperWhy this matters: salt and sugar tell the whole story

Salt and sugar illustrate the broad contrast. Salt (NaCl, ionic) melts at 801 °C, and salt water conducts electricity because dissolved Na⁠+ and Cl⁠- ions carry charge. Sugar (sucrose, molecular) melts and starts to decompose near 186 °C, and a sugar solution has low conductivity because sucrose dissolves as neutral molecules. Similar-looking crystals can therefore have very different structures and properties.

Go deeperCommon mistake: no metal does not always mean molecular

The composition shortcut has exceptions in both directions. NH⁠4Cl contains no metal, yet it is ionic because it consists of NH⁠4⁠+ and Cl⁠- ions. Anhydrous AlBr⁠3 contains a metal and a nonmetal, yet it is covalent and commonly forms Al⁠2Br⁠6 molecules. SiO⁠2 contains covalent bonds but forms an extended network rather than discrete molecules.

Use composition to form a first hypothesis, then ask whether the formula represents ions, discrete molecules, or a network solid.

Naming Binary Ionic Compounds

Binary ionic compounds contain exactly two elements: a metal cation and a nonmetal anion. The naming rules are straightforward:

  1. Name the cation first (simply the element name): sodium, calcium, aluminum
  2. Name the anion second, changing the element's ending to -ide: chlorine → chloride, oxygen → oxide, sulfur → sulfide

Examples:

  • NaCl = sodium chloride
  • CaO = calcium oxide
  • Al⁠2O⁠3 = aluminum oxide
  • MgBr⁠2 = magnesium bromide

Notice that the formula subscripts are not included in the name for ionic compounds — they are implied by the charges. This is a key difference from molecular compound naming.

Go deeperWhy this matters: one syllable on a label

Potassium chloride (KCl) is sold in grocery stores as a salt substitute. Potassium chlorate (KClO⁠3) is the oxidizer in match heads. One syllable separates a seasoning from a fire starter, which is why chemical nomenclature is deliberately rigid: the name must encode the exact composition, unambiguously, for anyone reading a reagent bottle, a shipping manifest, or an ingredient list. The -ide ending you are learning here is a promise: just the two elements, nothing more.

Writing Formulas from Names

To write a formula from a compound's name, you need to know the charges on the ions and then balance them so the compound is electrically neutral.

The process:

  1. Identify the cation and anion from the name
  2. Determine each ion's charge (from periodic table position or memorization)
  3. Find the simplest ratio that balances the total positive and negative charges

For example, “calcium chloride”: Ca⁠2+ and Cl⁠-. To balance 2+ with the 1− charges, you need two chloride ions: CaCl⁠2. Another example: “aluminum oxide”: Al⁠3+ and O⁠2-. The lowest common multiple of 3 and 2 is 6, requiring 2 Al and 3 O: Al⁠2O⁠3.

A quick shortcut: cross the charge numbers as subscripts (reduce to lowest ratio). Ca⁠2+ + Cl⁠- → CaCl⁠2. Al⁠3+ + O⁠2- → Al⁠2O⁠3.

Go deeperCommon mistake: writing 1:1 because the name sounds 1:1

“Calcium chloride” names two things, so students write CaCl. But ionic names deliberately omit counts: the subscripts are your job, reconstructed from the charges. Ca⁠2+ needs two Cl⁠- to balance, so calcium chloride is CaCl⁠2, and nothing in the name says so explicitly.

This is the sharpest difference from molecular naming, where dinitrogen tetroxide spells out its subscripts. Ionic names trust the charges; molecular names spell the counts. Mixing up the two systems (expecting “calcium dichloride”) is a sign the compound class was misidentified in step one.

Naming Compounds with Polyatomic Ions

Many ionic compounds contain polyatomic ions — charged groups of covalently bonded atoms that behave as a single unit. When naming these compounds, treat the polyatomic ion as a single entity and use its established name.

Examples:

  • NaNO⁠3 = sodium nitrate
  • Ca(OH)⁠2 = calcium hydroxide
  • NH⁠4Cl = ammonium chloride
  • (NH⁠4)⁠2SO⁠4 = ammonium sulfate

When a formula requires more than one polyatomic ion, place parentheses around the ion and add the subscript outside: Ca(NO⁠3)⁠2 means one Ca⁠2+ and two NO⁠3⁠- ions. The parentheses are essential — without them, the subscript applies only to the last atom.

Go deeperCommon mistake: breaking the ion apart

Within a formula such as Ca(NO⁠3)⁠2, treat nitrate as a named unit. The parentheses show that the subscript 2 multiplies the entire NO⁠3 group, giving two N atoms and six O atoms; CaNO⁠32 is incorrect notation because its subscripts do not represent two nitrate ions.

Polyatomic ions often remain intact when they act as spectator ions, although they can also participate in reactions. Keeping the ion together while naming and counting atoms prevents errors in later equation-balancing and net-ionic work.

Go deeperWhy this matters: polyatomic names on the medicine shelf

Milk of magnesia is magnesium hydroxide, Mg(OH)⁠2. Baking soda is sodium bicarbonate, NaHCO⁠3. Tums is calcium carbonate, CaCO⁠3. The polyatomic-ion vocabulary you are memorizing is the ingredient language of the pharmacy aisle, and reading it correctly means knowing instantly that all three products deliver a different metal attached to a different (or differently protonated) ion. The names are not trivia: hydroxide, bicarbonate, and carbonate neutralize stomach acid at different strengths and speeds.

Naming Binary Molecular Compounds

Binary molecular compounds (two nonmetals) use Greek prefixes to indicate the number of each atom, unlike ionic compounds which rely on charge balance.

The prefixes:

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

Rules:

  1. Name the first element with its prefix (omit mono- for the first element)
  2. Name the second element with its prefix and -ide ending
  3. Use the conventional vowel elisions found before oxide: monooxide becomes monoxide, tetraoxide becomes tetroxide, pentaoxide becomes pentoxide, and decaoxide becomes decoxide

Examples: N⁠2O⁠4 = dinitrogen tetroxide; PCl⁠5 = phosphorus pentachloride; CO = carbon monoxide; CO⁠2 = carbon dioxide.

Go deeperWhy this matters: mono- versus di- can be life or death

Carbon dioxide (CO⁠2) is what you exhale with every breath. Carbon monoxide (CO) binds to hemoglobin about 200 times more tightly than oxygen does and kills quietly in unventilated rooms; it is why homes have CO detectors. One Greek prefix carries the entire distinction.

Nitrogen runs the same lesson three times: dinitrogen monoxide (N⁠2O) is the dentist’s laughing gas, nitrogen monoxide (NO) is a signaling molecule your own blood vessels make, and nitrogen dioxide (NO⁠2) is a toxic brown component of smog. Same two elements, three different compounds, and only the prefixes keep them apart.

Naming Acids

Acids produce H⁠3O⁠+ ions in water, and acid names refer to the appropriate aqueous context. Their naming depends on whether the anion contains oxygen:

Binary acids (no oxygen, formula H + nonmetal):

  • Use the pattern: hydro- + root of element + -ic acid
  • HCl(aq) = hydrochloric acid; HBr(aq) = hydrobromic acid; H⁠2S(aq) = hydrosulfuric acid. Outside aqueous acid context, HCl is hydrogen chloride.

Oxyacids (contain oxygen, formula H + polyatomic ion with O):

  • If the anion ends in -ate, the acid ends in -ic acid: SO⁠4⁠2- (sulfate) → H⁠2SO⁠4 = sulfuric acid
  • If the anion ends in -ite, the acid ends in -ous acid: SO⁠3⁠2- (sulfite) → H⁠2SO⁠3 = sulfurous acid

The key pattern for the common oxyacids in this course is -ate → -ic acid and -ite → -ous acid.

Go deeperA memory hook that names dozens of acids

The suffix machine is small but productive. Remember one silly sentence: “I ate something icky” (-ate ions give -ic acids), and its partner “brite lights are dangerous” (-ite ions give -ous acids). With those two hooks plus the common polyatomic ions, you can generate sulfuric, sulfurous, nitric, nitrous, phosphoric, chloric, and chlorous acid without a table.

The hydrochloric acid pattern (hydro- + -ic for oxygen-free acids) also names the acid present in gastric juice.

Hydrates

A hydrate is an ionic compound that includes a specific number of water molecules within its crystal structure. The water is written after a centered dot in the formula, and named with a Greek prefix + “hydrate.”

Examples:

  • CuSO⁠4·5H⁠2O = copper(II) sulfate pentahydrate (blue crystals)
  • Na⁠2CO⁠3·10H⁠2O = sodium carbonate decahydrate (washing soda)
  • CoCl⁠2·6H⁠2O = cobalt(II) chloride hexahydrate

When a hydrate is heated, the water is driven off and the compound becomes anhydrous (without water). This often causes a visible color change — for example, hydrated cobalt(II) chloride is pink, while the anhydrous form is blue. Cobalt-chloride materials have traditionally been used as humidity indicators, although many current products use cobalt-free indicators.

Go deeperWhy this matters: hydrates in the hardware store

Plaster of Paris sets into a rigid cast because of hydrate chemistry: heating gypsum (CaSO⁠4·2H⁠2O) drives off most of its water; mixing the resulting powder with water lets the crystal take its water back and re-form solid gypsum. Traditional cobalt-chloride indicators illustrate the reverse color response: the material is blue when anhydrous and turns pink as it takes on water.

The dot in CuSO⁠4·5H⁠2O is real bookkeeping: those five waters have mass and occupy positions in the crystal structure. Losing them changes the substance’s appearance and formula mass.

Predicting Ionic Charges from the Periodic Table

The periodic table helps predict the common simple monatomic ions used in introductory chemistry:

  • Group 1 metals form 1+ cations (Li⁠+, Na⁠+, K⁠+)
  • Group 2 metals form 2+ cations (Mg⁠2+, Ca⁠2+, Ba⁠2+)
  • Aluminum commonly forms 3+ (Al⁠3+)
  • Common halide ions carry 1− (F⁠-, Cl⁠-, Br⁠-, I⁠-)
  • Common Group 16 anions include O⁠2- and S⁠2-
  • Common Group 15 anions include N⁠3- and P⁠3-

For these examples, metals lose electrons and nonmetals gain electrons to reach a noble-gas electron count. The charge equals the number of electrons lost or gained.

This pattern is scoped to common simple ions; it is not a universal oxidation-state rule for every element in a group. Many transition metals have more than one common charge (Fe⁠2+ and Fe⁠3+, Cu⁠+ and Cu⁠2+), so their ionic names use Roman numerals to state the charge.

Go deeperTest yourself: four charges from table position alone

Without looking anything up, predict the common simple ion for each element: barium, aluminum, sulfur, bromine. Then check below.

Answers: Ba⁠2+ (Group 2 loses two), Al⁠3+ (loses three), S⁠2- (gains two to reach argon’s electron count), Br⁠- (gains one to reach krypton’s electron count). This nearest-noble-gas reasoning works for these common ions. For elements with several accessible oxidation states, use chemical context rather than this shortcut alone.

Writing Formulas Using Charge Balance

An ionic compound is electrically neutral overall: the total positive charge equals the total negative charge. This charge balance principle determines the formula subscripts.

Step-by-step method:

  1. Write the symbols and charges of the cation and anion
  2. Find the least common multiple (LCM) of the charge magnitudes
  3. Divide the LCM by each ion's charge to get the subscripts
  4. Write the formula: cation first, anion second, subscripts as needed

Example: iron(III) oxide. Fe⁠3+ and O⁠2-. LCM of 3 and 2 = 6. Subscript for Fe = 6/3 = 2. Subscript for O = 6/2 = 3. Formula: Fe⁠2O⁠3.

For polyatomic ions, enclose the ion in parentheses when an outside subscript is greater than 1: Ca⁠2+ and NO⁠3⁠- → Ca(NO⁠3)⁠2. When nitrate appears once, write AgNO3; Ag(NO3) is unnecessary notation.

Diagnostic for a proposed introductory ionic formula. Check that (1) the stated cation and anion charges balance, (2) the chosen charges are the ones specified or expected in the problem, and (3) the subscripts are the smallest whole-number ratio. Thus FeNO3 would require Fe⁠+, which is outside the common Fe⁠2+/Fe⁠3+ ions used here, while Cu3SO4 would require a nonintegral average charge per equivalent Cu site in this simple model. A formula that fails these checks does not represent the stated ions in the expected introductory model.

Go deeperCommon mistake: forgetting to reduce the ratio

Balancing tin(IV) with oxide by brute force, some students write Sn⁠2O⁠4: two 4+ charges against four 2− charges does balance. But ionic formulas report the smallest whole-number ratio, so the answer is SnO⁠2. An ionic crystal is not made of Sn⁠2O⁠4 molecules; it is an endless lattice, and the formula only records the ratio of ions in it.

Run the reduction check every time both subscripts are even (or share any common factor). The LCM method avoids the trap automatically: LCM of 4 and 2 is 4, giving 4/4 = 1 tin and 4/2 = 2 oxygens directly.

Roman Numerals for Transition Metal Compounds

Many transition metals (and some main-group metals like Pb and Sn) form common ions with different charges. To specify which charge is present, a Roman numeral in parentheses follows the metal name.

Examples:

  • FeCl⁠2 = iron(II) chloride (Fe⁠2+)
  • FeCl⁠3 = iron(III) chloride (Fe⁠3+)
  • CuO = copper(II) oxide (Cu⁠2+)
  • Cu⁠2O = copper(I) oxide (Cu⁠+)

To determine the Roman numeral from a formula, calculate the metal charge from the known anion charges. In FeCl⁠3, three Cl⁠- ions contribute 3−, so Fe is 3+ and the name is iron(III) chloride.

Metals treated as fixed-charge in introductory nomenclature, including Group 1, Group 2, Al⁠3+, Zn⁠2+, and Ag⁠+, are normally named without Roman numerals in this course.

Go deeperWhy this matters: the numeral on your iron supplement

Many iron supplements use ferrous salts because ferrous forms are generally more bioavailable than comparable ferric forms. Corrosion products called rust can include several iron(III) oxides and oxyhydroxides, such as hematite and goethite. The iron charge therefore carries useful information about composition and behavior. Pharmacists, water-treatment engineers, and materials chemists all rely on that distinction.

Go deeperCommon mistake: the numeral is the charge, not the count

Iron(III) chloride happens to be FeCl⁠3, so students sometimes conclude the III counts chlorines. It does not: it states the charge on the metal ion. Iron(III) oxide is Fe⁠2O⁠3, not FeO⁠3; the III still means Fe⁠3+, while the subscripts come from charge balance.

The reverse slip is writing sodium(I) chloride. Metals treated as fixed-charge in introductory nomenclature (Group 1, Group 2, Al⁠3+, Zn⁠2+, and Ag⁠+) are normally named without a numeral in this course. Use the numeral when the ionic name needs to identify which relevant metal charge is present.

Common Polyatomic Ions

A relatively small set of polyatomic ions appears repeatedly in chemistry. Memorizing them is essential for naming and writing formulas:

  • Positive: NH⁠4⁠+ (ammonium)
  • −1 charge: OH⁠- (hydroxide), NO⁠3⁠- (nitrate), NO⁠2⁠- (nitrite), ClO⁠3⁠- (chlorate), C⁠2H⁠3O⁠2⁠- (acetate), MnO⁠4⁠- (permanganate), HCO⁠3⁠- (bicarbonate), CN⁠- (cyanide)
  • −2 charge: SO⁠4⁠2- (sulfate), SO⁠3⁠2- (sulfite), CO⁠3⁠2- (carbonate), CrO⁠4⁠2- (chromate), Cr⁠2O⁠7⁠2- (dichromate), HPO⁠4⁠2- (hydrogen phosphate)
  • −3 charge: PO⁠4⁠3- (phosphate)

Notice a helpful naming pattern: many polyatomic ions come in pairs. The -ate form has more oxygen atoms; the -ite form has one fewer oxygen. Nitrate (NO⁠3⁠-) vs. nitrite (NO⁠2⁠-). Sulfate (SO⁠4⁠2-) vs. sulfite (SO⁠3⁠2-). The charge stays the same between pairs.

Go deeperThe oxygen ladder shrinks the memorization

The chlorine oxyanions form a four-rung ladder that one memorized ion unlocks: perchlorate (ClO⁠4⁠-, one O more than -ate), chlorate (ClO⁠3⁠-, the reference -ate form), chlorite (ClO⁠2⁠-, one O less), and hypochlorite (ClO⁠-, two O less). Learn chlorate and the prefixes per- and hypo- generate the rest; all four carry the same 1− charge, because only the oxygen count changes.

One rung is in your laundry room: household bleach is sodium hypochlorite, NaClO. The same ladder logic extends to bromine and iodine oxyanions, so a single pattern covers a dozen ions.

The -ous and -ic Naming Conventions

An older (but still widely used) naming system uses the suffixes -ous and -ic to distinguish between lower and higher oxidation states of metals:

  • -ous = lower charge: ferrous (Fe⁠2+), cuprous (Cu⁠+), stannous (Sn⁠2+)
  • -ic = higher charge: ferric (Fe⁠3+), cupric (Cu⁠2+), stannic (Sn⁠4+)

These older names often use Latin roots for the metal: iron → ferr-, copper → cupr-, tin → stann-, lead → plumb-.

The same -ous/-ic pattern appears in acid naming: the -ite polyatomic ion gives an -ous acid, and the -ate ion gives an -ic acid. Sulfite → sulfurous acid; sulfate → sulfuric acid.

While the modern IUPAC system using Roman numerals is preferred, you will encounter -ous/-ic names in older references and on many reagent bottles.

Go deeperWhy this matters: the old names live on store shelves

Check a toothpaste box and you may find “stannous fluoride” (SnF⁠2, tin(II) fluoride); check an iron supplement and you will likely find “ferrous sulfate” (iron(II) sulfate). The -ous/-ic system predates IUPAC nomenclature by generations, and consumer products, pharmacology, and older literature never let it go. You are learning both systems not for historical completeness but because both are in active use: the modern one in your coursework and journals, the Latin one on the products in your bathroom.

Classifying Compounds by Structure

For the introductory naming set in LO 4.1, first decide whether the formula represents ions, discrete molecules, or a network-covalent solid.

  • Ionic compounds consist of cations and anions. Name the cation first and the anion second; monatomic anions take an -ide ending, while polyatomic ions keep their established names. Example: CuSO4 = copper(II) sulfate.
  • Molecular compounds consist of discrete covalent molecules and commonly contain only nonmetals. Binary molecular names use Greek prefixes. Example: N2O5 = dinitrogen pentoxide.
  • Network-covalent solids, such as SiO⁠2, contain an extended covalent network rather than discrete molecules.

Composition is a useful first clue, but it is not an absolute classifier: NH⁠4Cl is ionic without a metal, and anhydrous AlBr⁠3 is covalent despite containing one. Formulas representing acids in water use the acid-naming route in LO 4.6; for example, HCl is hydrogen chloride outside aqueous acid context and hydrochloric acid in water.

Nomenclature Decision Framework and Common Mistakes

Naming compounds correctly requires a systematic decision path:

  1. Identify the structural class (LO 4.1): does the formula represent ions, discrete molecules, or a network-covalent solid? Use composition as a first clue and remember the stated exceptions.
  2. If ionic: identify the cation and anion. If the metal has more than one relevant charge, use a Roman numeral (LO 4.10). Name the cation first, then the anion; a monatomic anion takes an -ide ending.
  3. If ionic with polyatomic ions: use the established polyatomic-ion name (LO 4.4). Watch for -ate versus -ite endings.
  4. If a binary molecular compound: use Greek prefixes for the atom counts and change the second element’s ending to -ide (LO 4.5). Omit mono- on the first element.
  5. If the formula is being named as an acid in water: apply the acid convention (LO 4.6). HX(aq) uses hydro-___-ic acid; an -ate oxyanion gives an -ic acid, while an -ite oxyanion gives an -ous acid.

Common mistakes include treating the composition heuristic as an absolute rule, omitting a needed Stock numeral, using molecular prefixes in ionic names, confusing -ate/-ite pairs, failing to balance ionic charges, and omitting the water-count prefix from a hydrate name.

Learning Objectives

After studying this topic, you should be able to:

  1. Classify a compound as ionic, molecular, or network covalent from its formula and context
  2. Name binary ionic compounds from their formulas
  3. Write formulas of ionic compounds from their names
  4. Name compounds containing polyatomic ions
  5. Name binary molecular (covalent) compounds using numerical prefixes
  6. Name acids (binary and oxyacids) from their formulas
  7. Identify and name hydrates
  8. Predict common ionic charges for main-group elements from periodic table position
  9. Write formulas for ionic compounds using charge balance
  10. Use the Stock system (Roman numerals) to name transition-metal compounds

How-To Procedure

How to Name a Chemical Compound from Its Formula

  1. Classify the compound (LO 4.1): identify whether it consists of ions, discrete molecules, or a network-covalent solid. Use metal/nonmetal composition as a first clue, then check for known exceptions such as NH4Cl, anhydrous AlBr3, and SiO2.
  2. For ionic compounds: name the cation (metal name), then the anion with an -ide ending for monatomic anions or the polyatomic ion's name as-is.
  3. If the metal is a transition metal or can have multiple charges, determine its charge from the anion charges and include a Roman numeral in parentheses.
  4. For molecular compounds that are not acids: use Greek prefixes on both elements (omit mono- on the first element) and change the second element's ending to -ide.
  5. For formulas named as acids in water: apply the acid-naming convention from LO 4.6. Binary acid (HX in water) → hydro-___-ic acid. Oxyacid from an -ate ion → ___-ic acid. Oxyacid from an -ite ion → ___-ous acid.
  6. Double-check: verify the formula is electrically neutral for ionic compounds, and confirm you used the correct prefix or Roman numeral.

Worked Example

Naming an Ionic Compound with a Transition Metal

Problem

A compound has the formula Cr⁠2(SO⁠4)⁠3. What is its systematic name?

Solution
  1. Identify the ions: Cr is a transition metal (cation); SO⁠4 is the sulfate polyatomic ion with a 2− charge
  2. Determine the charge on chromium: the compound is neutral, so total positive charge = total negative charge
  3. Three sulfate ions contribute 3 × (2−) = 6− total negative charge
  4. Two chromium ions must contribute 6+ total, so each Cr is 3+ → Cr⁠3+
  5. Name the cation with Roman numeral: chromium(III)
  6. Name the anion: sulfate
  7. Combine: chromium(III) sulfate
Answer

Cr⁠2(SO⁠4)⁠3 is chromium(III) sulfate.

Test Your Understanding

A student names the compound FeCl2 as 'iron chloride.' What is wrong with this name, and what information is missing? How would the name differ for FeCl3?

Practice Problems

conceptual

Classify each compound as ionic or molecular and provide its systematic name; for molecular compounds that are also acids, apply the acid-naming convention rather than the Greek-prefix one: (a) MgBr2, (b) N2O5, (c) HClO3 (acid), (d) Cu2SO4.

calculation

Write the chemical formula for each compound: (a) calcium phosphate, (b) dinitrogen monoxide, (c) iron(III) hydroxide, (d) sulfurous acid.

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Self-Study Questions

What is the difference between an ionic compound and a molecular (covalent) compound?

How do you name a binary ionic compound from its formula?

When are Roman numerals needed in the name of an ionic compound?

Hint: Think about which metals can form more than one type of cation.

What is a polyatomic ion? Name three common examples.

How do you write the formula of an ionic compound from its name?

How do you name a binary molecular compound using prefixes?

What is the naming convention for acids that do not contain oxygen (binary acids)?

What is the naming convention for oxyacids?

What is a hydrate and how is it named?

How do you determine the charge on a monatomic ion from the periodic table?

What is the difference between the -ous and -ic suffixes in older naming systems?

How do you classify a compound as ionic, molecular, or network covalent from its formula and context?

Hint: Use the element types as a first clue, then decide whether the substance contains ions, discrete molecules, or an extended covalent network.

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Related Content

Study Topics

Atomic StructureReactions in Solution

Elements

NaClFeCuS

Reference Tables

Common Ions

Content Sources

Concept sections adapted from open educational resources under Creative Commons licensing:

  • OpenStax Chemistry 2e, Ch 2.6: Ionic and Molecular Compounds (CC BY 4.0)
  • OpenStax Chemistry 2e, Ch 2.7: Chemical Nomenclature (CC BY 4.0)
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