Chapter 2 Notes: 2.6-2.8 — Molecules, Ions, and Nomenclature

2.6 Molecules and Molecular Compounds

• Quick goal for section 2.6-2.8: learn naming and classification of ionic vs molecular (covalent) compounds; memorize seven diatomic elements; be able to classify a given formula/name as ionic or molecular.

• Ionic compounds (general):

  • Composed of ions held together by strong electrostatic forces in a crystal lattice; typically form extended, repeating lattice structures, not discrete molecules.

  • In chemical formulas, you see elemental symbols with subscripts that indicate the simplest whole-number ratio of ions (empirical formula) required for charge neutrality (example: $ext{NaCl}$ is a 1:1 $ext{Na}^+$ to $ext{Cl}^-$ ratio).

  • When dissolved in water (or in molten state), their ions dissociate, making them good electrical conductors.

  • Often involve a metal (which forms a cation) and a nonmetal (which forms an anion) or a polyatomic ion. The strong electrostatic attraction between these oppositely charged ions forms the ionic bond.

  • The common type of ionic compound is metal + nonmetal; can also be metal + polyatomic ion (e.g., $ext{NH}<em>4 ext{Cl}$ where $ext{NH}</em>4^+$ is a polyatomic cation).

• Molecular compounds (covalent):

  • Composed of discrete molecules, each containing a specific number of atoms bonded covalently; they do not form extended lattices in the same way ionic compounds do.

  • Typically formed from two or more nonmetals sharing electrons to achieve stable electron configurations.

  • A molecular formula gives the actual number of atoms of each element in one molecule (e.g., $ext{H}<em>2 ext{O}$, $ext{CO}</em>2$).

  • Most molecular compounds contain only nonmetals; bonds inside molecules are covalent, involving shared electron pairs.

  • Unlike ionic compounds, molecular compounds generally do not conduct electricity when dissolved in water, as they do not dissociate into ions (unless they are acids, which is a special category).

  • To understand bonding in these compounds more deeply, see chapter 8 (bonding concepts).

• Molecular vs empirical formulas:

  • Molecular formula: tells the actual number of atoms of each element in one molecule of the compound. It provides the true composition of a single molecule. $ext{e.g., } ext{C}<em>6 ext{H}</em>{12} ext{O}_6$ (glucose), $ext{H}<em>2 ext{O}</em>2$ (hydrogen peroxide).

  • Empirical formula: the simplest, lowest whole-number ratio of atoms in a compound. For ionic compounds, the empirical formula is always used because they don't form discrete molecules but rather repeating formula units in a lattice. For molecular compounds, the empirical formula can be derived if only the ratio is known, or if it represents the simplest whole-number ratio.

  • Examples:

    • Glucose has molecular formula $ext{C}<em>6 ext{H}</em>{12} ext{O}<em>6$; its empirical formula is $ext{CH}</em>2 ext{O}$ (derived by dividing all subscripts by 6, the greatest common divisor).

    • Formaldehyde has molecular formula $ext{CH}<em>2 ext{O}$; its empirical formula is also $ext{CH}</em>2 ext{O}$ (molecular and empirical formulas can be the same).

    • Acetic acid has molecular formula $ext{C}<em>2 ext{H}</em>4 ext{O}<em>2$; its empirical formula is $ext{CH}</em>2 ext{O}$.

    • Water: molecular formula $ext{H}<em>2 ext{O}$; empirical formula is also $ext{H}</em>2 ext{O}$ (the ratio is already 1:2:1).

    • Hydrogen peroxide: molecular formula $ext{H}<em>2 ext{O}</em>2$; empirical formula is $ext{HO}$ (divide by 2).

    • Benzene: molecular formula $ext{C}<em>6 ext{H}</em>6$; empirical formula $ext{CH}$ (divide by 6).

• From formulas to molecular visuals:

  • The molecular formula shows the actual identity and number of atoms in one molecule. Some molecules are elements (e.g., $ext{H}<em>2, ext{O}</em>2, ext{N}<em>2$) and some are compounds (e.g., $ext{H}</em>2 ext{O}, ext{CO}*2$).

  • Structural formulas (e.g., Lewis structures, developed in chapter 8) indicate how atoms are connected within a molecule, often showing specific bonds (single, double, triple).

  • Diatomic molecules (elements that exist naturally as diatomic molecules): memorize seven species because they are often encountered in chemical reactions as such.

• The seven diatomic molecules (memorization aid):

  • These elements exist naturally as molecules composed of two atoms chemically bonded together, rather than as individual atoms.

  • The mnemonic: “Have No Fear of Ice Cold Beer” corresponds to:

    • $ext{H}*2$ (Hydrogen)

    • $ext{N}*2$ (Nitrogen)

    • $ext{F}*2$ (Fluorine)

    • $ext{O}*2$ (Oxygen)

    • $ext{I}*2$ (Iodine)

    • $ext{Cl}*2$ (Chlorine)

    • $ext{Br}*2$ (Bromine)

  • Important note: While Hydrogen is placed at the left edge of the periodic table, it is a nonmetal. The other six (N, F, O, I, Cl, Br) form a '7' shape on the periodic table (starting at Nitrogen and going across to Fluorine, then down to Iodine).

• Practical implications:

  • When you see a nonmetal–nonmetal combination in a compound, expect a molecular compound with covalent bonds.

  • When you see a metal–nonmetal or polyatomic ion combination, expect an ionic compound with a lattice structure.

• Quick example walkthroughs (molecular perspective):

  • $ext{CH}*4$ (methane): consists of a central carbon atom covalently bonded to four hydrogen atoms. Its formula is a molecular formula, not an empirical formula, showing it's a discrete molecule.

  • Structural formulas use lines to represent covalent bonds (e.g., H-C-H where lines denote single bonds). 3D models (ball-and-stick, space-filling) help visualize molecular geometry, bond angles, and electron cloud distribution.

• Summary points to remember for 2.6:

  • Ionic compounds are extended crystal lattices formed by electrostatic attraction between ions; molecular compounds are composed of discrete molecules formed by covalent bonds.

  • Molecular formula gives actual atom counts in one molecule; empirical formula gives the lowest whole-number ratio of atoms.

  • Seven diatomic molecules ($ext{H}<em>2, ext{N}</em>2, ext{F}<em>2, ext{O}</em>2, ext{Cl}<em>2, ext{Br}</em>2, ext{I}*2$) must be memorized as they are elemental species existing in diatomic form.

  • Diatomic molecules are elements, not compounds, when they are in their pure elemental form.

2.7 Ions and Ionic Compounds

• Atoms vs ions:

  • Ions are formed by the gain or loss of electrons from a neutral atom, resulting in a net electrical charge. They are not formed by changes in the nucleus (protons or neutrons).

  • Loss of one or more electrons yields a positively charged ion (cation); this typically occurs with metals which have lower ionization energies.

  • Gain of one or more electrons yields a negatively charged ion (anion); this typically occurs with nonmetals which have higher electron affinities.

  • Ionization is the process of forming ions, often to achieve a stable electron configuration, typically resembling a noble gas.

• Ion notation:

  • Ions are written as elemental symbols with a superscript indicating the charge in the top-right corner. The magnitude of the charge is written before the sign (e.g., $ext{Na}^+$ for a +1 charge, $ext{O}^{2-}$ for a -2 charge).

  • Examples: $ext{Fe}^{2+}$ (iron(II) ion), $ext{Fe}^{3+}$ (iron(III) ion), $ext{Cl}^-$.

  • Monatomic ions are ions consisting of a single atom (e.g., $ext{Na}^+$, $ext{Mg}^{2+}$, $ext{Al}^{3+}$).

• Cations vs anions:

  • Cations (positive ions) are typically formed from metals (Groups 1A, 2A, 3A, and transition metals) that readily lose valence electrons.

  • Anions (negative ions) are typically formed from nonmetals (Groups 5A, 6A, 7A) that readily gain electrons into their valence shell.

  • Hydrogen can form a cation (H$^+$, often called a proton in aqueous solution, common in acids) and can form a hydride anion (H$^-$) when bonded to a highly electropositive metal (e.g., $ext{NaH}$).

  • A neutral atom has an equal number of protons and electrons and is written without a superscript.

• Common monatomic ions and patterns to memorize (high level):

  • Group 1A metals (alkali metals: Li, Na, K, Rb, Cs): always form $+1$ ions (e.g., $ext{Na}^+$).

  • Group 2A metals (alkaline earth metals: Be, Mg, Ca, Sr, Ba): always form $+2$ ions (e.g., $ext{Mg}^{2+}$).

  • Group 3A (e.g., Al): Aluminum always forms a $+3$ ion (e.g., $ext{Al}^{3+}$).

  • Group 5A nonmetals: Nitrogen typically forms a $-3$ ion as a common monatomic ion (nitride, $ext{N}^{3-}$, in ionic contexts).

  • Group 6A nonmetals (chalcogens: O, S, Se, Te): typically form $-2$ ions (e.g., oxide $ext{O}^{2-}$, sulfide $ext{S}^{2-}$).

  • Group 7A nonmetals (halogens: F, Cl, Br, I): typically form $-1$ ions (e.g., chloride $ext{Cl}^-$, bromide $ext{Br}^-$, iodide $ext{I}^-$, fluoride $ext{F}^-$, especially when not bonded to oxygen or another halogen).

  • Hydrogen: typically $+1$ as H$^+$ (proton), but $-1$ in metal hydrides (H$^-$).

  • Transition metals: often have variable charges (e.g., $ext{Fe}^{2+}$, $ext{Fe}^{3+}$, $ext{Cu}^+$, $ext{Cu}^{2+}$). Their charge must be indicated using Roman numerals in their names.

  • Three special transition metals with fixed charges to memorize: These do not require Roman numerals.

    • $ext{Ag}^+$ (silver ion, $+1$ charge)

    • $ext{Zn}^{2+}$ (zinc ion, $+2$ charge)

    • $ext{Cd}^{2+}$ (cadmium ion, $+2$ charge)

• Ion notation and ion class names:

  • Cations: The element name is followed by the word “ion” (e.g., $ext{Na}^+$ is sodium ion, $ext{Mg}^{2+}$ is magnesium ion, $ext{Al}^{3+}$ is aluminum ion; for transition metals, e.g., $ext{Fe}^{2+}$ is iron(II) ion).

  • Monatomic Anions: The stem of the element name is used, and the ending is changed to -ide (e.g., $ext{O}^{2-}$ is oxide, $ext{N}^{3-}$ is nitride, $ext{Cl}^-$ is chloride, $ext{S}^{2-}$ is sulfide).

• Balancing ionic compounds (example problem-solving approach):

  • Ionic compounds must be electrically neutral, meaning the total positive charge from cations must equal the total negative charge from anions.

  • To determine the charge on a transition metal in a compound, assign the known charges to the other ion(s) and then determine the charge needed for the transition metal to balance the compound to zero net charge.

  • Example: In $ext{CuO}$, oxygen is known to form an oxide ion, $ext{O}^{2-}$. To achieve a neutral compound (total charge = 0), copper must have a $+2$ charge. So, it is copper(II) oxide. ($ext{Cu}^{2+}$ and $ext{O}^{2-}$).

  • Example: In $ext{Cu}*2 ext{O}$, oxygen is still $ext{O}^{2-}$. With two copper atoms, each copper must contribute $+1$ charge to balance the $ext{O}^{2-}$. Thus, it is copper(I) oxide. (2 x $ext{Cu}^+$ and $ext{O}^{2-}$).

• Oxidation numbers (oxidation state) as a bookkeeping tool:

  • Oxidation numbers (or oxidation states) are hypothetical charges assigned to atoms in a molecule or ion if all bonds were purely ionic. They help in tracking electrons and balancing chemical equations (especially redox reactions).

  • Rule set (summary):

    • Rule 1: The oxidation number of an atom in a pure element (uncombined with other elements) is $0$ (e.g., $ext{N}<em>2, ext{O}</em>2, ext{Fe}, ext{Na}$ all have an oxidation number of $0$).

    • Rule 2: The oxidation number of a monatomic ion is equal to its charge (e.g., $ext{Na}^+ = +1$, $ext{Zn}^{2+} = +2$, $ext{Cl}^- = -1$).

    • Rule 3 (nonmetals in compounds): Assign hydrogen, oxygen, and halogens with their common oxidation states based on electronegativity, typically:

      • Hydrogen: usually $+1$ when bonded to nonmetals (e.g., in $ext{H}*2 ext{O}, ext{HCl}$). It is $-1$ when bonded to metals to form hydrides (e.g., in $ext{NaH}$).

      • Oxygen: usually $-2$ in most compounds (e.g., $ext{CO}<em>2, ext{SO}</em>2$). Exceptions: in peroxides (e.g., $ext{H}<em>2 ext{O}</em>2, ext{Na}<em>2 ext{O}</em>2$), each oxygen is $-1$. In superoxides (e.g., $ext{KO}_2$), each oxygen is $-1/2$. When bonded to fluorine (e.g., $ext{OF}*2$), oxygen is $+2$ because fluorine is more electronegative.

      • Halogens (Group 7A): typically $-1$ in most compounds, especially binary compounds with less electronegative elements (e.g., $ext{NaCl}, ext{HBr}$). However, they can have positive oxidation states when bonded to oxygen or more electronegative halogens (e.g., in oxyanions like $ext{ClO}<em>4^-$ where Cl is $+7$; in $ext{BrF}</em>5$ where Br is $+5$).

    • Rule 4: The sum of the oxidation numbers for all atoms in a neutral compound must be zero. For a polyatomic ion, the sum of the oxidation numbers must equal the charge of the ion.

  • Example of applying oxidation numbers: For $ext{PO}*4^{3-}$ (phosphate ion):

    • Oxygen is usually $-2$. So, for four oxygen atoms, the total contribution is $4 imes (-2) = -8$.

    • The overall charge of the phosphate ion is $-3$.

    • Let P's oxidation state be $x$. Then $x + (-8) = -3$. Solving for $x$, P must have an oxidation state of $+5$.

    • Oxidation states: $ext{P}^{+5}$ and $ext{O}^{-2}$ (four of them).

• Polyatomic ions (summary):

  • Polyatomic ions are groups of two or more atoms covalently bonded together that possess an overall net positive or negative charge. They behave as a single unit when forming ionic compounds.

  • Ammonium ( ext{NH}*4^+}) is the most common positively charged polyatomic ion (cation).

  • Most polyatomic ions are negatively charged (anions), and their names do not end in -ide (with the exception of hydroxide $( ext{OH}^-)$, cyanide $( ext{CN}^-)$, and peroxide $( ext{O}*2^{2-})$. These are common enough to remember).

  • Examples: sulfate ($ext{SO}<em>4^{2-}$), nitrate ($ext{NO}</em>3^-$, phosphate ($ext{PO}<em>4^{3-}$, carbonate ($ext{CO}</em>3^{2-}$, hydroxide ($ext{OH}^-$).

  • Naming rule of oxyanions (polyatomic ions containing oxygen, a subset of polyatomic ions):

    • The suffixes -ate and -ite indicate the relative number of oxygen atoms for the same central nonmetal atom.

      • The -ate ending typically denotes the form with more oxygen atoms (e.g., $ext{nitrate } ext{NO}<em>3^-$, $ext{sulfate } ext{SO}</em>4^{2-}$).

      • The -ite ending typically denotes the form with fewer oxygen atoms (e.g., $ext{nitrite } ext{NO}<em>2^-$, $ext{sulfite } ext{SO}</em>3^{2-}$).

    • For oxyanions with four possible oxygen states, prefixes per- and hypo- are used:

      • The per- prefix (e.g., perchlorate $ext{ClO}*4^-$) indicates one more oxygen atom than the -ate form.

      • The hypo- prefix (e.g., hypochlorite $ext{ClO}^-$, for one less oxygen than chlorite) indicates one fewer oxygen atom than the -ite form.

  • Relationship between oxyanions and oxyacids (covered in 2.8):

    • If the oxyanion name ends in -ate, the corresponding acid name ends in -ic acid (e.g., nitrate $ext{NO}<em>3^-$ forms nitric acid $ext{HNO}</em>3$; sulfate $ext{SO}<em>4^{2-}$ forms sulfuric acid $ext{H}</em>2 ext{SO}*4$).

    • If the oxyanion name ends in -ite, the corresponding acid name ends in -ous acid (e.g., nitrite $ext{NO}<em>2^-$ forms nitrous acid $ext{HNO}</em>2$; sulfite $ext{SO}<em>3^{2-}$ forms sulfurous acid $ext{H}</em>2 ext{SO}*3$).

• Polyatomic ion practice and examples:

  • Ammonium sulfate: Cation is $ext{NH}<em>4^+$ and anion is $ext{SO}</em>4^{2-}$. To balance charges, two ammonium ions are needed for every one sulfate ion. Formula: $( ext{NH}<em>4)</em>2 ext{SO}*4$.

  • Potassium phosphate: Cation is $ext{K}^+$ and anion is $ext{PO}<em>4^{3-}$. To balance, three potassium ions for every one phosphate. Formula: $ext{K}</em>3 ext{PO}*4$.

  • Magnesium nitride: Cation is $ext{Mg}^{2+}$ and anion is $ext{N}^{3-}$. Using the criss-cross method (explained below): $ext{Mg}<em>3 ext{N}</em>2$.

• Criss-cross method (swap-and-drop) for ionic formulas:

  • This is a useful shortcut for determining the subscripts in an ionic compound to ensure charge neutrality.

  • Write the ions side by side with their charges. The numerical value of the cation's charge becomes the subscript for the anion, and the numerical value of the anion's charge becomes the subscript for the cation.

  • Reduce the subscripts to the lowest whole-number ratio (if possible).

  • If a polyatomic ion requires a subscript greater than 1, enclose the polyatomic ion in parentheses before applying the subscript.

  • Example: For $ext{Mg}^{2+}$ and $ext{NO}<em>3^-$. The 2 from $ext{Mg}^{2+}$ becomes the subscript for $ext{NO}</em>3^-$; the 1 from $ext{NO}<em>3^-$ becomes the subscript for $ext{Mg}^{2+}$. Formula: $ext{Mg}( ext{NO}</em>3)*2$ (parentheses are essential for the polyatomic ion).

  • Example: For $ext{Al}^{3+}$ and $ext{O}^{2-}$. The 3 from $ext{Al}^{3+}$ becomes the subscript for $ext{O}^{2-}$; the 2 from $ext{O}^{2-}$ becomes the subscript for $ext{Al}^{3+}$. Formula: $ext{Al}<em>2 ext{O}</em>3$.

• Summary points for 2.7:

  • Ionic compounds are formed by electron transfer (metals lose to form cations, nonmetals gain to form anions); polyatomic ions are groups of atoms with a charge that behave as single units.

  • Common monatomic ions have predictable charges based on their periodic table group; transition metals often have variable charges, requiring Roman numerals in their names (except fixed-charge $ext{Ag}^+, ext{Zn}^{2+}, ext{Cd}^{2+}$).

  • Oxidation numbers are a formal tool for assigning charges to atoms within compounds or ions, following a specific set of rules, which helps in naming and understanding chemical reactivity.

  • For naming and writing formulas, it is crucial to memorize common ions (both monoatomic and polyatomic); use the criss-cross method (and parentheses for polyatomic ions) to balance charges for neutral ionic compounds.

2.8 Nomenclature of Inorganic Compounds

• Three major naming categories (nomenclature):

  • Binary ionic compounds: composed of a metal and a nonmetal, forming simple ions.

  • Compounds with polyatomic ions: involving polyatomic ions as either the cation or anion.

  • Binary molecular compounds: composed of two different nonmetals.

  • Acids: a special category related to hydrogen and either a nonmetal or an oxyanion.

• IUPAC system and memorization prerequisites:

  • IUPAC stands for the International Union of Pure and Applied Chemistry; this is the globally recognized system chemists use to standardize names unambiguously, ensuring clear communication.

  • To efficiently name compounds and write formulas, you must memorize common monatomic ions, variable charges of transition metals (and their exceptions), and a list of common polyatomic ions (charges and names).

• Binary ionic compounds (name-the-ions approach):

  • Name the cation first, followed by the anion.

  • For monoatomic cations from Group 1A, 2A, and Al, use the element's name directly (e.g., sodium, magnesium, aluminum).

  • For monoatomic anions, take the stem of the element's name and change the ending to -ide (e.g., $ext{S}^{2-}$ becomes sulfide; $ext{O}^{2-}$ becomes oxide; $ext{Cl}^-$ becomes chloride).

  • For transition metals (and some main group metals like Pb, Sn) that can have variable charges, use Roman numerals in parentheses immediately after the metal name to indicate its specific charge (e.g., $ext{FeCl}<em>2$ is iron(II) chloride because $ext{Cl}</em>2$ represents 2 x $-1$, so Fe must be $+2$; $ext{FeCl}<em>3$ is iron(III) chloride because $ext{Cl}</em>3$ represents 3 x $-1$, so Fe must be $+3$).

  • For the three special transition metals with fixed charges ($ext{Ag}^+, ext{Zn}^{2+}, ext{Cd}^{2+}$), no Roman numeral is used in their names because their charge is invariant (e.g., $ext{AgCl}$ is silver chloride, not silver(I) chloride; $ext{ZnCl}*2$ is zinc chloride).

  • Examples:

    • $ext{Na}*2 ext{S}$ → sodium sulfide (Na is Group 1A, fixed $+1$ charge; S is a monoatomic anion $-2$).

    • $ext{MgCl}*2$ → magnesium chloride (Mg is Group 2A, fixed $+2$ charge; Cl is a monoatomic anion $-1$).

    • $ext{Al}<em>2 ext{O}</em>3$ → aluminum oxide (Al is Group 3A, fixed $+3$ charge; O is a monoatomic anion $-2$).

• Polyatomic ions in ionic naming:

  • When polyatomic ions are present, treat the entire polyatomic ion as a single unit; do not change its name or break it apart in the overall compound name.

  • If the polyatomic ion is the cation, it is named first (e.g., ammonium ($ext{NH}*4^+$)).

  • If the polyatomic ion is the anion, it is named second, using its memorized name (e.g., nitrate ($ext{NO}<em>3^-$, sulfate ($ext{SO}</em>4^{2-}$, phosphate ( ext{PO}*4^{3-}$)).

  • Examples: ext{NaNO}3$→ sodium nitrate;$ ext{K}3 ext{PO}4$→ potassium phosphate;$( ext{NH}4)3 ext{PO}4$would be ammonium phosphate (parentheses are used in the formula to denote multiple polyatomic ions when needed).</p></li></ul></li><li><p>Criss-cross method for ionic formulas (summary):</p><ul><li><p>This method helps directly derive the formula from the charges of the ions. Write down the cation and anion with their charges. Swap the numerical values of the charges to become the subscripts. Reduce the subscripts to the lowest whole-number ratio if possible. Use parentheses for polyatomic ions needing a subscript greater than one.</p></li><li><p>Example: For$ ext{Mg}^{2+}$and chloride ($ ext{Cl}^-$). Swap the '2' from Mg and the '1' from Cl. This results in$ ext{MgCl}*2$. (The 1 for Mg is typically omitted).</p></li><li><p>Example: For$ ext{Fe}^{3+}$and sulfate ($ ext{SO}4^{2-}$). Swap the '3' from Fe and the '2' from sulfate. This results in$ ext{Fe}2( ext{SO}4)3$. Parentheses are crucial for sulfate.</p></li></ul></li><li><p>Special cases in naming and charges:</p><ul><li><p>The three transition metals with fixed charges (Ag ($+1$), Zn ($+2$), Cd ($+2$)) are unique and are named without Roman numerals in the IUPAC system. This is a crucial memorization point.</p></li><li><p>All other transition metals (and some post-transition metals like Tin and Lead) use Roman numerals to clearly indicate their specific oxidation state (charge) when forming ions with variable charges.</p></li></ul></li><li><p>Binary molecular compounds (two nonmetals):</p><ul><li><p>The name is constructed similarly to ionic naming but uses numerical prefixes because covalent bonds lead to discrete molecules with varying numbers of atoms.</p></li><li><p>Use prefixes to indicate the number of atoms of <em>each</em> element in the molecule. Prefixes: mono- (1), di- (2), tri- (3), tetra- (4), penta- (5), hexa- (6), hepta- (7), octa- (8), nona- (9), deca- (10).</p></li><li><p><strong>Crucial Rule</strong>: The prefix <strong>mono-</strong> is <strong>never used for the <em>first</em> element</strong> in the name if there is only one atom of that element. It <em>is</em> used for the second element if there is only one atom of it.</p></li><li><p>The second element’s name always ends in <strong>-ide</strong>.</p></li><li><p><strong>Important spelling rule</strong>: When the prefix ends in 'a' or 'o' (like tetra-, penta-, mono-) and the element name begins with a vowel (e.g., oxide), the 'a' or 'o' from the prefix is often dropped for easier pronunciation (e.g., carbon monoxide, not carbon monooxide; dinitrogen pentoxide, not dinitrogen pentaoxide).</p></li></ul></li><li><p>Common examples of binary molecular naming:</p><ul><li><p>$ ext{CO}$→ carbon monoxide (mono- is used for oxygen, but not for carbon since it's the first element).</p></li><li><p>$ ext{CO}*2$→ carbon dioxide.</p></li><li><p>$ ext{N}2 ext{O}5$→ dinitrogen pentoxide (penta- becomes pent- before 'oxide').</p></li><li><p>$ ext{SO}*3$→ sulfur trioxide.</p></li><li><p>$ ext{CCL}*4$→ carbon tetrachloride.</p></li><li><p>Dihydrogen monoxide ($ ext{H}*2 ext{O}$) is the systematic name for water, though rarely used in practice, it's a good illustration of the rules.</p></li></ul></li><li><p>Acids (special category):</p><ul><li><p>Acids are compounds that produce$ ext{H}^+$ions when dissolved in water. Their naming depends on whether they are binary acids or oxyacids.</p></li><li><p><strong>Binary acids</strong> (hydrogen + a nonmetal, typically Group 7A or S):</p><ul><li><p>The name starts with the prefix <strong>hydro-</strong>.</p></li><li><p>The stem of the nonmetal's name follows.</p></li><li><p>The suffix <strong>-ic acid</strong> is added.</p></li><li><p>Examples:$ ext{HCl}$→ <strong>hydro</strong>chlor<strong>ic acid</strong>;$ ext{HBr}$→ <strong>hydro</strong>brom<strong>ic acid</strong>;$ ext{H}*2 ext{S}$→ <strong>hydro</strong>sulfur<strong>ic acid</strong>.</p></li></ul></li><li><p><strong>Oxyacids</strong> (acids containing hydrogen, oxygen, and another nonmetal; derived from polyatomic oxyanions): no <strong>hydro-</strong> prefix is used.</p><ul><li><p>The name of the acid is directly based on the polyatomic oxyanion from which it is derived:</p><ul><li><p>If the oxyanion ends in <strong>-ate</strong>, the acid name ends in <strong>-ic acid</strong> (e.g., nitrate ($ ext{NO}3^-$) → nitric acid ($ ext{HNO}3$); sulfate ($ ext{SO}4^{2-}$) → sulfuric acid ($ ext{H}2 ext{SO}*4$)).</p></li><li><p>If the oxyanion ends in <strong>-ite</strong>, the acid name ends in <strong>-ous acid</strong> (e.g., nitrite ($ ext{NO}2^-$) → nitrous acid ($ ext{HNO}2$); sulfite ($ ext{SO}3^{2-}$) → sulfurous acid ($ ext{H}2 ext{SO}*3$)).</p></li></ul></li><li><p>For oxyacids with four possible oxygen states (like chlorine):</p><ul><li><p><strong>Per-</strong> prefix (more oxygen than -ate): perchlorate ($ ext{ClO}4^-$) → perchloric acid ($ ext{HClO}4$).</p></li><li><p><strong>Hypo-</strong> prefix (less oxygen than -ite): hypochlorite ($ ext{ClO}^-$) → hypochlorous acid ($ext{HClO}$).