All topics summary

8.1 Molecular Compounds

Molecules and Molecular Compounds

• Only noble gas elements like helium and neon exist as uncombined atoms (monatomic).

• Not all elements are monatomic; some, like oxygen (O2), exist as diatomic molecules, where two atoms are bonded together.

• Ionic compounds are typically crystalline solids with high melting points.

• Molecular compounds can have different properties:

  • Water (H2O) is liquid at room temperature.

  • Carbon dioxide (CO2) and nitrous oxide (N2O) are gases at room temperature.

• Ionic bonds involve atoms giving up or accepting electrons, while covalent bonds involve atoms sharing electrons.

• Atoms held together by sharing electrons are joined by a covalent bond.

• A molecule is a neutral group of atoms joined by covalent bonds.

• Oxygen gas consists of oxygen molecules, each with two covalently bonded oxygen atoms.

• Diatomic molecules contain two atoms (e.g., hydrogen, nitrogen, halogens).

• A molecular compound is composed of molecules (e.g., water).

Representing Molecules

• A molecular formula is the chemical formula of a molecular compound.

• It shows how many atoms of each element are in a substance.

• The subscript indicates the number of atoms of each element; the subscript 1 is omitted.

• Butane's molecular formula is C4H10, indicating four carbon atoms and ten hydrogen atoms.

• Subscripts are not necessarily the lowest whole-number ratios; molecular formulas reflect the actual number of atoms.

• Oxygen molecule (O2) consists of two oxygen atoms bonded together.

• Molecular formulas do not show a molecule's structure or which atoms are covalently bonded to one another.

• Molecular structure refers to the arrangement of atoms within a molecule.

• Carbon dioxide's molecular structure shows the three atoms arranged in a row, with the carbon atom in the middle between the two oxygen atoms.

Comparing Molecular and Ionic Compounds

• The representative unit of a molecular compound is a molecule.

• The representative unit of an ionic compound is a formula unit, which is the lowest whole-number ratio of ions.

• Molecules are made up of two or more atoms acting as a unit.

• No discrete units exist in ionic compounds; they consist of a continuous array of ions.

• There is no such thing as a molecule of sodium chloride or magnesium chloride.

• Molecular compounds tend to be gases or liquids at room temperature, while ionic compounds are solids.

• Molecular compounds are typically composed of atoms of two or more nonmetals, whereas ionic compounds are formed from a metal and a nonmetal.

• Molecular compounds tend to have lower melting and boiling points than ionic compounds.

Key Concepts

• Covalent bond: a bond formed by the sharing of electrons between atoms.

• Molecule: a neutral group of atoms joined by covalent bonds.

• Diatomic molecule: a molecule consisting of two atoms.

• Molecular compound: a compound composed of molecules.

• Molecular formula: a chemical formula of a molecular compound that shows the kinds and numbers of atoms present in a molecule of a compound.

BIG IDEA

• In molecular compounds, bonding occurs when atoms share electrons.

• In ionic compounds, bonding occurs when electrons are transferred between atoms.

The Octet Rule in Covalent Bonding

• In covalent bonds, electron sharing usually occurs so that atoms attain the electron configurations of noble gases.

  • A single hydrogen atom has one electron, but two hydrogen atoms share electrons to form a covalent bond in a diatomic hydrogen molecule.

  • Each hydrogen atom attains the electron configuration of helium, a noble gas with two electrons.

• Combinations of atoms of the nonmetals and metalloids in Groups 4A, 5A, 6A, and 7A are likely to form covalent bonds.

• The combined atoms usually acquire a total of eight electrons, or an octet, by sharing electrons, so that the octet rule applies.

• The hydrogen atoms in a hydrogen molecule are held together mainly by the attraction of the shared electrons to the positive nuclei.

• Two atoms held together by sharing one pair of electrons are joined by a single covalent bond.

• Hydrogen gas consists of diatomic molecules whose atoms share only one pair of electrons, forming a single covalent bond.

• An electron dot structure such as H:H represents the shared pair of electrons of the covalent bond by two dots.

• The pair of shared electrons forming the covalent bond is also often represented as a dash, as in H—H for hydrogen.

• A structural formula represents the covalent bonds as dashes and shows the arrangement of covalently bonded atoms.

• The halogens also form single covalent bonds in their diatomic molecules. Fluorine is one example.

• By sharing electrons and forming a single covalent bond, two fluorine atoms each achieve the electron configuration of neon.

• In the F2 molecule, each fluorine atom contributes one electron to complete the octet.

• A pair of valence electrons that is not shared between atoms is called an unshared pair, also known as a lone pair or a nonbinding pair.

• In F2, each fluorine atom has three unshared pairs of electrons.

• The oxygen atom in water has two unshared pairs of valence electrons.

• Methane contains four single covalent bonds.

• The carbon atom has four valence electrons and needs four more valence electrons to attain a noble-gas configuration.

• Each of the four hydrogen atoms contributes one electron to share with the carbon atom, forming four identical carbon–hydrogen bonds.

• When carbon forms bonds with other atoms, it usually forms four bonds, as in methane.

• The formation of four bonds by carbon can be explained by the fact that one of carbon’s 2s electrons is promoted to the vacant 2p orbital to form the following electron configuration.

• This electron promotion requires only a small amount of energy, and the stability of the resulting methane more than compensates for the small energy cost.

Sample Problem 8.1 Drawing an Electron Dot Structure

Hydrochloric acid (HCl (aq)) is prepared by dissolving gaseous hydrogen chloride (HCl (g)) in water. Hydrogen chloride is a diatomic molecule with a single covalent bond. Draw the electron dot structure for HCl.

• In a single covalent bond, a hydrogen and a chlorine atom must share a pair of electrons. Each must contribute one electron to the bond.

• Draw the electron dot structures for the hydrogen and chlorine atoms. Through electron sharing, the hydrogen and chlorine atoms attain the electron configurations of the noble gases helium and argon, respectively.

Double and Triple Covalent Bonds

• Atoms form double or triple covalent bonds if they can attain a noble gas structure by sharing two or three pairs of electrons.

• A double covalent bond is a bond that involves two shared pairs of electrons.

• Similarly, a bond formed by sharing three pairs of electrons is a triple covalent bond.

• The carbon dioxide (CO2) molecule contains two oxygens, each of which shares two electrons with carbon to form a total of two carbon–oxygen double bonds.

• Nitrogen (N2), a major component of Earth’s atmosphere, contains triple bonds.

• A single nitrogen atom has five valence electrons; each nitrogen atom in the molecule must share three electrons to have the electron configuration of neon.

• You might think that an oxygen atom, with six valence electrons, would form a double bond by sharing two of its electrons with another oxygen atom.

• In such an arrangement, all the electrons within the molecule would be paired.

• Experimental evidence, however, indicates that two of the electrons in O2 are still unpaired.

• Thus, the bonding in the oxygen molecule (O2) does not obey the octet rule.

Diatomic Elements

• The “octet” in the octet rule refers to eight of what?

• Each of the atoms joined by a covalent bond usually acquires eight electrons in its valence shell.

• Most noble gases have eight valence electrons.

Coordinate Covalent Bonds

• How are coordinate covalent bonds different from other covalent bonds?

• Carbon monoxide (CO) is an example of a type of covalent bonding different from that seen in water, ammonia, methane, and carbon dioxide.

• It is possible for both carbon (which needs to gain four electrons) and oxygen (which needs to gain two electrons) to achieve noble-gas electron configurations by a type of bonding called coordinate covalent bonding.

• Look at the double covalent bond between carbon and oxygen. With the double bond in place, the oxygen had a stable electron configuration, but the carbon does not.

• As shown below, the dilemma is solved if the oxygen also donates one of its unshared pairs of electrons for bonding.

• A covalent bond in which one atom contributes both bonding electrons is a coordinate covalent bond.

• In a coordinate covalent bond, the shared electron pair comes from one of the bonding atoms.

• The ammonium ion (NH4+) consists of atoms joined by covalent bonds, including a coordinate covalent bond.

• A polyatomic ion, such as NH4+, is a tightly bound group of atoms that has a positive or negative charge and behaves as a unit.

• The ammonium ion forms when a positively charged hydrogen ion (H+) attaches to the unshared electron pair of an ammonia molecule (NH3).

• Most polyatomic cations and anions contain covalent and coordinate covalent bonds.

• Therefore, compounds containing polyatomic ions include both ionic and covalent bonding.

Sample Problem 8.2 Drawing the Electron Dot Structure of a Polyatomic Ion

The H3O+ ion forms when a hydrogen ion is attracted to an unshared electron pair in a water molecule. Draw the electron dot structure of the hydronium ion.

• Each atom must share electrons to satisfy the octet rule.

• Draw the electron dot structure of the water molecule and the hydrogen ion. Then draw the electron dot structure of the hydronium ion.

• The oxygen must share a pair of electrons with the added hydrogen ion to form a coordinate covalent bond.

• Check that all the atoms have the electrons they need and that the charge is correct. The oxygen in the hydronium ion has eight valence electrons, and each hydrogen shares two valence electrons, satisfying the octet rule. The water molecule is neutral, and the hydrogen ion has a positive charge, giving the hydronium ion a charge of 1+.

• Do all atoms joined in covalent bonds donate electrons to the bond?

• No. In coordinate covalent bonds, the shared electron pair comes from one of the bonding atoms.

Exceptions to the Octet Rule

• What are some exceptions to the octet rule?

• The octet rule cannot be satisfied in molecules whose total number of valence electrons is an odd number. There are also molecules in which an atom has less, or more, than a complete octet of valence electrons.

• Two plausible electron dot structures can be drawn for the NO2 molecule, which has a total of seventeen valence electrons.

• It is impossible to draw an electron dot structure for NO2 that satisfies the octet rule for all atoms, yet NO2 does exist as a stable molecule.

• Some molecules with an even number of valence electrons, such as some compounds of boron, also fail to follow the octet rule.

• A few atoms, especially phosphorus and sulfur, expand the octet to ten or twelve electrons. Sulfur hexafluoride (SF6) is an example.

• Are molecules that do not obey the octet rule necessarily unstable?

• No. There are molecules like NO2 that do not obey the octet rule, but that are stable, naturally occurring molecules.

Bond Dissociation Energies

• How is the strength of a covalent bond related to its bond dissociation energy?

• A large quantity of heat is released when hydrogen atoms combine to form hydrogen molecules.

• This release of heat suggests that the product is more stable than the reactants.

• The covalent bond in the hydrogen molecule (H2) is so strong that it would take 435 kJ of energy to break apart all of the bonds in 1 mole (about 2 grams) of H2.

• The energy required to break the bond between two covalently bonded atoms is known as the bond dissociation energy.

• The units for this energy are often given in kJ/mol, which is the energy needed to break one mole of bonds.

• A large bond dissociation energy corresponds to a strong covalent bond.

• A typical carbon–carbon single bond has a bond dissociation energy of 347 kJ/mol.

• Strong carbon–carbon bonds help explain the stability of carbon compounds; they are unreactive partly because the dissociation energy is high.

• True or False: A strong covalent bond has a low bond dissociation energy.

• False. A large bond dissociation energy corresponds to a strong covalent bond.

Resonance

• How are resonance structures used?

• The ozone molecule has two possible electron dot structures.

• Notice that the structure on the left can be converted to the one on the right by shifting electron pairs without changing the positions of the oxygen atoms.

• The two electron dot structures for ozone are examples of what are still referred to as resonance structures.

• Resonance structures are structures that occur when it is possible to draw two or more valid electron dot structures that have the same number of electron pairs for a molecule or ion.

Key Concepts

• In covalent bonds, electron sharing occurs so that atoms attain the configurations of noble gases.

• In a coordinate covalent bond, the shared electron pair comes from a single atom.

• The octet rule is not satisfied in molecules with an odd number of electrons and in molecules in which an atom has less, or more, than a complete octet of valence electrons.

• A large bond dissociation energy corresponds to a strong covalent bond.

• In ozone, the bonding of oxygen atoms is a hybrid of the extremes represented by the resonance forms.

Glossary Terms

• Single covalent bond: a bond formed when two atoms share a pair of electrons

• Structural formula: a chemical formula that shows the arrangement of atoms in a molecule or a polyatomic ion; each dash between a pair of atoms indicates a pair of shared electrons

• Unshared pair: a pair of valence electrons that is not shared between atoms

• Double covalent bond: a bond in which two atoms share two pairs of electrons

• Triple covalent bond: a covalent bond in which three pairs of electrons are shared by two atoms

• Coordinate covalent bond: a covalent bond in which one atom contributes both bonding electrons

• Polyatomic ion: a tightly bound group of atoms that behaves as a unit and has a positive or negative charge

• Bond dissociation energy: the energy required to break the bond between two covalently bonded atoms; this value is usually expressed in kJ per mol of substance

• Resonance structure: one of the two or more equally valid electron dot structures of a molecule or polyatomic ion

9. 1 Naming Ions

Monatomic Ions

• How can you determine the charges of monatomic ions?

• Ionic compounds consist of a positive metal ion and a negative nonmetal ion combined in a proportion such that their charges add up to a net charge of zero.

• For example, the ionic compound sodium chloride (NaCl) consists of one sodium ion (Na+) and one chloride ion (Cl–).

• It is important, in learning the language of chemistry, to be able to name and write the chemical formulas for all ionic compounds.

• The first step is to learn about the ions that form ionic compounds.

• Some ions, called monatomic ions, consist of a single atom with a positive or negative charge resulting from the loss or gain of one or more valence electrons, respectively.

• Recall that metallic elements tend to lose valence electrons.

Cations

• All the Group 1A ions have a 1+ charge (Li+, Na+ , K+, Rb+, and Cs+).

• Group 2A metals, including magnesium and calcium, tend to lose two electrons to form cations with a 2+ charge (Mg2+ and Ca2+).

• Aluminum is the only common Group 3A metal and tends to lose three electrons to form a 3+ cation (Al3+).

• When the metals in Groups 1A, 2A, and 3A lose electrons, they form cations with positive charges equal to their group number.

• The names of the cations of Group 1A, Group 2A, and Group 3A metals are the same as the name of the metal, followed by the word ion or cation.

• Thus, Na+ is the sodium ion (or cation), Ca2+ is the calcium ion (or cation), and Al3+ is the aluminum ion (or cation).

Anions

• Nonmetals tend to gain electrons to form anions, so the charge of a nonmetallic ion is negative.

• The charge of any ion of a Group A nonmetal is determined by subtracting 8 from the group number.

• The elements in Group 7A form anions with a 1– charge (7 – 8 = –1).

• Anion names start with the stem of the element name and end in -ide.

• For example, two elements in Group 7A are fluorine and chlorine. The anions for these nonmetals are the fluoride ion (F–) and the chloride ion (Cl–).

• Anions of nonmetals in Group 6A have a 2– charge (6 – 8 = –2).

• Group 6A elements, oxygen and sulfur, form the oxide anion (O2–) and the sulfide anion (S2–), respectively.

• The first three elements in Group 5A, nitrogen, phosphorus, and arsenic, can form anions with a 3– charge (5 – 8 = –3).

• These anions have the symbols N3– , P3– , and As3– and are called, respectively, nitride ion, phosphide ion, and arsenide ion.

Metals That Form More Than One Ion

• Many of the transition metals (Groups 1B–8B) form more than one cation with different ionic charges.

• The charges of the cations of many transition metal ions must be determined from the number of electrons lost.

• For example, the transition metal iron forms two common cations, Fe2+ (two electrons lost) and Fe3+ (three electrons lost).

• Cations of tin and lead, the two metals in Group 4A, can also have more than one common ionic charge.

• Two methods are used to name ions that can have more than one common ionic charge. The preferred method is called the Stock system.

• In the Stock system, you place a Roman numeral in parentheses after the name of the element to indicate the numerical value of the charge.

• For example, the cation Fe2+ is named iron(II) ion and is read “iron two ion.” No space is left between the element name and the Roman numeral in parentheses.

• The Fe3+ ion is named iron(III) ion and is read “iron three ion.”

• An older, less useful method for naming these cations uses a root word with different suffixes at the end of the word.

• The older, or classical, name of the element is used to form the root name for the element.

• For example, ferrum is Latin for iron, so ferr- is the root name for iron.

• The suffix -ous is used to name the cation with the lower of the two ionic charges.

• The suffix -ic is used with the higher of the two ionic charges. Using this system, Fe2+ is the ferrous ion, and Fe3+ is the ferric ion.

• You can usually identify an element from what may be an unfamiliar classical name by looking for the element’s symbol in the name.

• For example, ferrous (Fe) is iron, cuprous (Cu) is copper, and stannous (Sn) is tin.

• A major disadvantage of using classical names for ions is that they do not tell you the actual charge of the ion.

Symbols and Names of Common Metal Ions With More Than One Ionic Charge

• A few transition metals have only one ionic charge.

• The names of these cations do not have a Roman numeral.

• These exceptions include silver, with cations that have a 1+ charge (Ag+), as well as cadmium and zinc, with cations that have a 2+ charge (Cd2+ and Zn2+).

Sample Problem 9.1 Naming Cations and Anions

Name the ion formed by each of the following elements:
a. potassium
b. lead, 4 electrons lost
c. sulfur

• You can use the periodic table to determine the charge of most Group A elements. Ions with positive charges are cations; ions with negative charges are anions. The names of nonmetallic anions end in -ide. Metallic cations take the name of the metal. Some metals, including transition metals, can form more than one cation. Use a Roman number in the Stock name or use the classical name with a suffix to name these metals.
  a. Following the rules for naming metallic cations, K+ is named potassium ion.
  b. Following the rules for naming metals that can form more than one cation, Pb4+ is named lead(IV) or plumbic ion.
  c. Following the rules for naming nonmetallic anions, S2– is named sulfide ion.

• What type of elements (metals or nonmetals) tends to form cations? What type of elements tends to form anions?

• Metals tend to form cations. Nonmetals tend to form anions.

CHEMISTRY & YOU

• For cations, the word ion or cation follows the name of the element.

• Metals that form more than one cation are named by adding a Roman numeral in parentheses to indicate the value of the charge after the name of the element, followed by the word ion.

• Anion names start with the stem of the element name and end in -ide.

Polyatomic Ions

• How do polyatomic ions differ from monatomic ions? How are they similar?

• Unlike a monatomic ion, a polyatomic ion is composed of more than one atom. But like a monatomic ion, a polyatomic ion behaves as a unit and carries a charge.

• The sulfate anion consists of one sulfur atom and four oxygen atoms.

• These five atoms together comprise a single anion with an overall 2– charge. The formula is written SO42–.

• The names and formulas of some common polyatomic ions are shown here. Note that the names of most polyatomic ions end in - ite or -ate.

• Sometimes the same two or three elements combine in different ratios to form different polyatomic ions. Look for pairs of ions for which there is both an -ite and an -ate ending, for example, sulfite and sulfate.

Common Polyatomic Ions

• Note the number of oxygen atoms and the endings on each name. You should be able to discern a pattern in the naming convention.

• The charge is the same on each polyatomic ion in a pair for which there is both an -ite and an -ate ion.

• The -ite ending indicates one less oxygen atom than the -ate ending.

• However, the ending does not tell you the actual number of oxygen atoms in the ion.

• For example, the nitrite ion has two oxygen atoms, and the sulfite ion has three oxygen atoms.

• When the formula for a polyatomic ion begins with H (hydrogen), you can think of the H as representing a hydrogen ion (H+) combined with another polyatomic ion.

• For example, HCO3 – is a combination of H+ and CO32–. Note that the charge on the new ion is the algebraic sum of the ionic charges of the two component ions.

Key Concepts

• When the metals in Groups 1A, 2A, and 3A lose electrons, they form cations with positive charges equal to their group number.

• The charge of any ion of a Group A nonmetal is determined by subtracting 8 from the group number.

• The charges of the cations of many transition metal ions must be determined from the number of electrons lost.

• Unlike a monatomic ion, a polyatomic ion is composed of more than one atom. But like a monatomic ion, a polyatomic ion behaves as a unit and carries a charge.

Glossary Terms

• Monatomic ion: a single atom with a positive or negative charge resulting from the loss or gain of one or more valence electrons

BIG IDEA

• An element’s position in the periodic table supplies information on ion formation and bonding tendencies, which is used to write the names and formulas of ions and compounds.

9. 2 Naming and Writing Formulas for Ionic Compounds

Binary Ionic Compounds

• How do you determine the formula and name of a binary ionic compound?

• Before the science of chemistry developed, compounds were often named to describe some property of the substance or its source.

• For example, a common name for potassium carbonate (K2CO3) is potash because the compound was obtained by boiling wood ashes in iron pots.

• NaHCO3 is called baking soda because it is used in baking to make baked goods rise.

Writing Formulas for Binary Ionic Compounds

• A binary compound is composed of two elements; they can be ionic or molecular.

• If you know the name of a binary ionic compound, you can write the formula by first writing the symbol of the cation and then the anion.

• Add subscripts as needed to balance the charges, ensuring the net ionic charge is zero.

• For example, potassium chloride (KCl) consists of potassium cations (K+) and chloride anions (Cl–) in a 1:1 ratio.

• Iron(III) oxide (Fe2O3) contains Fe3+ cations and O2– anions; two Fe3+ ions balance three O2– ions.

Crisscross Method

• Another approach is the crisscross method, where the numerical value of each ion's charge becomes the subscript for the other ion, dropping the signs.

• If Ca2+ and S2– are combined, the formula Ca2S2 is obtained; the 2:2 ratio is not the lowest whole-number ratio.

• The correct formula for calcium sulfide is CaS.

Sample Problem 9.2 Writing Formulas for Binary Ionic Compounds

Write the formulas for the following binary ionic compounds:
a. copper(II) sulfide
b. potassium nitride

The symbol for the cation appears first in the formula for the compound. The ionic charges in an ionic compound must balance, and the ions must be combined in the lowest whole-number ratio.

a. CuS: 1(2+) + 1(2–) = 0
b. K3N: 3(1+) + 1(3–) = 0

Naming Binary Ionic Compounds

• If you know the formula for a binary ionic compound, you can write its name, provided it consists of a monatomic metallic cation and a monatomic nonmetallic anion.

Naming Rules

• Place the cation name first, followed by the anion name. Cs2O is cesium oxide; NaBr is sodium bromide; SrF2 is strontium fluoride.

• If the metallic element has more than one common ionic charge, include a Roman numeral (e.g., copper(I) or copper(II)).

• CuO requires determining the copper cation's charge; since the oxide anion is always 2–, copper must be 2+ to balance the charge.

CHEMISTRY & YOU

• Many companies use sodium sulfite (Na2SO3) to keep dried fruit looking delicious.

• Is Na2SO3 a binary compound? Explain: Na2SO3 is not a binary compound because binary compounds are composed of two elements. SO3 is a compound, not an element.

Sample Problem 9.3 Naming Binary Ionic Compounds

Name the following binary ionic compounds:
a. CoI2
b. Li2Se

the cation name followed by the anion name. The name of a metal ion that has more than one common ionic charge must include a Roman numeral indicating the charge.

In the correct order name:
a. cobalt(II) iodide
b. lithium selenide
*

• Why is it necessary to balance the charges of the two ions in a binary ionic compound?

• A binary ionic compound carries no charge when the charges of the ions that combine to form it are balanced.

Compounds With Polyatomic Ions

• How do you determine the formula and name of a compound with a polyatomic ion?

• Seashells are made of calcium carbonate (CaCO3), which is not a binary compound because it contains more than two elements.

• An -ate or -ite ending indicates a polyatomic anion including oxygen.

Writing Formulas for Compounds With Polyatomic Ions

• Write the symbol/formula for the cation, then the anion, balancing the charges with subscripts.

• For calcium nitrate, Ca2+ and NO3– require two nitrate anions, hence Ca(NO3)2 (parentheses indicate multiple polyatomic ions).

• Whenever more than one polyatomic ion is needed to balance the charges in an ionic compound, use parentheses to set off the polyatomic ion in the formula.

• Lithium carbonate requires two lithium cations: Li2CO3.

Sample Problem 9.4 Writing Formulas for Compounds With Polyatomic Ions

What are the formulas for these ionic compounds?
a. magnesium hydroxide
b. potassium sulfate
the formula for each ion in the order listed in the name. Use subscripts to balance the charges.
i. If more than one polyatomic ion is needed to balance a formula, place the polyatomic ion formula in parentheses, followed by the appropriate subscript.

• magnesium hydroxide: Mg(OH)2

• potassium sulfate: K2SO4

Naming Compounds With Polyatomic Ions

• Identify polyatomic ions, state the cation name, then the anion name, including a Roman numeral for metallic elements with multiple charges.

• NaClO is sodium hypochlorite; the rules remain the same as those of binary naming.

Sample Problem 9.5 Naming Compounds With Polyatomic Ions

Name the following ionic compounds.
a. (NH4)2C2O4
b. Fe(ClO3)3
l is not a polyatomic ion: 1;to name the compound, list the names of the ions in the order written in the formula—the cation name followed by the anion name. The name of an ion that has more than one common ionic charge must include a Roman numeral indicating the charge. 1

• (NH4)2C2O4: ammonium oxalate

• Fe(ClO3)3: iron(III) chlorate
  *

• What is the difference between binary ionic compounds and compounds with polyatomic ions?

• Binary ionic compounds are made of two ions, each made of just one element. Compounds with polyatomic ions can contain ions made of just one element, but they also contain a polyatomic ion made of multiple elements.

Key Concepts

• To write the formula of a binary ionic compound, first write the symbol of the cation and then the anion. Then balance the charges.

• The name of a binary ionic compound is the cation name followed by the anion name.

• To write formulas for compounds with polyatomic ions, write the symbol for the cation followed by the symbol for the anion. Then balance the charges.

• To name a compound containing a polyatomic ion, state the cation name followed by the anion name.

Glossary Terms

• Binary compound: a compound composed of two elements; NaCl and Al2O3 are binary compounds

9. 3 Naming and Writing Formulas for Molecular Compounds

Binary Molecular Compounds

• What guidelines are used to write the name and formula of a binary molecular compound?

• Binary ionic compounds are composed of the ions of two elements, a metal and a nonmetal.

• Binary molecular compounds are composed of two nonmetals and are not ions.

• Binary molecular compounds are composed of molecules, not ions, so ionic charges cannot be used to write formulas or to name them.

Differences between CO and CO2

• Carbon and oxygen can combine to form CO and CO2, which have very different properties.

• You exhale CO2 as a product of your body chemistry, and it is normally present in the air you breathe. CO is a poisonous gas that interferes with your blood’s ability to carry oxygen to body cells.

Prefixes for Names

• The prefixes in the names of binary molecular compounds tell how many atoms of an element are present in each molecule of the compound.

• The prefix mono- would be used for the single oxygen atom in CO.

• The prefix di- would be used for the two oxygen atoms in CO2.

Prefixes Used in Naming Binary Molecular Compounds