Comprehensive Guide to Chemical Reactions and Chemical Nomenclature

Fundamentals of Chemicals and Matter

• A chemical is defined as either an element or a compound possessing a distinct, invariant set of properties.

• Properties of chemicals are classified into two primary categories:

  • Physical properties: Characteristics that can be directly observed or measured without altering the chemical identity of the substance. Key examples include melting point, boiling point, color, and density.

  • Chemical properties: Characteristics that can only be observed when a substance undergoes a chemical interaction or reaction with another chemical substance. Key examples include flammability (reaction with oxygen), solubility (the ability to dissolve in water), and reactivity with acids.

• Everything in the physical universe is composed of chemicals ("EVERYTHING is chemicals!").

• Labeling commercial products such as bottled water as "Chemical Free" is scientifically inaccurate, as water ($H_2O$) and all naturally dissolved minerals within it are entirely chemical substances.

Chemical Reactions and Equation Anatomy

• A chemical reaction occurs when chemical substances undergo a transformation into completely different chemical substances.

• Chemical equations represent reactions using standard structural components:

  • Reactants: The initial substances that react together and are consumed throughout the reaction. They are written on the left side of the reaction arrow.

  • Products: The new chemical substances formed by the conclusion of the reaction. They are written on the right side of the reaction arrow.

  • Reaction Arrow: Denoted by $\rightarrow$, pointing from reactants toward products to represent the direction of chemical change.

• Chemical notation conventions:

  • Species: The specific chemical element or compound involved.

  • Coefficient: The large full-sized number placed in front of a chemical species indicating how many equivalents or units of that substance react or are produced in a balanced chemical equation.

  • Subscript: The small number positioned to the lower right of an element symbol indicating the exact count of atoms of the element situated immediately to its left.

  • Superscript: Positioned at the upper right of a chemical species indicating the overall net electrical charge of that ion or polyatomic species.

  • Parentheses: Used to group polyatomic ions. A subscript positioned outside the parentheses multiplies every atom contained within.

  • Absence of a Number: In coefficients, subscripts, or superscripts, an unwritten number implicitly denotes a value of $1$.

• Quantitative Example (Determining Atom Count from Formulas):

  • For the compound $Fe_2(SO_4)_3$:

  • The exterior subscript $3$ indicates there are $3$ individual sulfate ($SO_4^{2-}$) ions present in the formula unit.

  • Each sulfate ion contains $4$ oxygen atoms.

  • Multiplying the internal subscript by the external subscript yields the total oxygen atom count:     $4 \times 3 = 12$

  • Therefore, $Fe_2(SO_4)_3$ contains exactly $12$ oxygen atoms.

Nomenclature of Simple Ionic Compounds

• General Nomenclature Principle: With the exception of water ($H_2O$), which must be memorized by its common name, all chemical compounds are named according to systematic, rule-based nomenclature.

• Ionic Compound Properties:

  • Formed by electrostatic attraction between positively charged cations and negatively charged anions held tightly together.

  • The net electrical charge of every ionic compound is strictly $0$.

  • Positive ($+$ ) and negative ($-$ ) charges must completely balance one another.

  • Ion charges are denoted using superscripts (e.g., $Ca^{2+}$ indicates a calcium cation with a $+2$ charge).

• Determining Charges from the Periodic Table:

  • Group I elements always form cations with a $+1$ charge.

  • Group II elements always form cations with a $+2$ charge.

  • Group III elements always form cations with a $+3$ charge.

  • Group V elements always form anions with a $-3$ charge.

  • Group VI elements always form anions with a $-2$ charge.

  • Group VII elements always form anions with a $-1$ charge.

• Charge Balancing and Formula Determination Examples:

  • $Na^+$ and $O^{2-}$ balance in a 2:1 ratio to form $Na_2O$

  • $Ca^{2+}$ and $Cl^-$ balance in a 1:2 ratio to form $CaCl_2$

  • $Al^{3+}$ and $S^{2-}$ balance in a 2:3 ratio to form $Al_2S_3$

  • $Mn^{7+}$ and $O^{2-}$ balance in a 2:7 ratio to form $Mn_2O_7$

  • $Mg^{2+}$ and $O^{2-}$ balance in a 1:1 ratio to form $MgO$

  • $Ca^{2+}$ and $N^{3-}$ balance in a 3:2 ratio to form $Ca_3N_2$

  • Calcium Fluoride: Combines $Ca^{2+}$ and $F^-$ to form $CaF_2$

• Rules for Converting Simple Ionic Formulas to Names (Non-Transition Metal Cations):

  • Cation: Name the metal element directly; do not include Roman numerals.

  • Monoatomic Anion: Name the nonmetal element, replacing its ending with the suffix "-ide".

  • Example: $NaCl$ is named sodium chloride.

  • Example: $K_2O$ is named potassium oxide.

  • Polyatomic Anion: Retain the standardized name of the polyatomic group without altering its ending.

Polyatomic Ions and Oxyanion Patterns

• A polyatomic ion is a charged molecular species composed of two or more atoms covalently bonded together that acts as a single discrete unit.

• Reference List of Common Polyatomic Ions:

  • Ammonium: $NH_4^+$ (the primary common positively charged polyatomic ion)

  • Acetate: $C_2H_3O_2^-$

  • Carbonate: $CO_3^{2-}$

  • Chlorate: $ClO_3^-$

  • Chlorite: $ClO_2^-$

  • Chromate: $CrO_4^{2-}$

  • Cyanide: $CN^-$

  • Dichromate: $Cr_2O_7^{2-}$

  • Hydrogen carbonate (bicarbonate): $HCO_3^-$

  • Hydrogen sulfate: $HSO_4^-$

  • Hydroxide: $OH^-$

  • Hypochlorite: $ClO^-$

  • Nitrate: $NO_3^-$

  • Nitrite: $NO_2^-$

  • Perchlorate: $ClO_4^-$

  • Permanganate: $MnO_4^-$

  • Peroxide: $O_2^{2-}$

  • Phosphate: $PO_4^{3-}$

  • Sulfate: $SO_4^{2-}$

  • Sulfite: $SO_3^{2-}$

  • Thiosulfate: $S_2O_3^{2-}$

• The "-ate" versus "-ite" Oxyanion Naming Rule:

  • Many polyatomic anions end in either "-ate" or "-ite".

  • Mnemonic: "-ate ate an extra oxygen atom."

  • The "-ate" form always has one more oxygen atom than the corresponding "-ite" form possessing the same base element and overall charge.

  • Paired Examples:

  • Sulfate is $SO_4^{2-}$, while Sulfite is $SO_3^{2-}$

  • Nitrate is $NO_3^-$, while Nitrite is $NO_2^-$

  • Chlorate is $ClO_3^-$, while Chlorite is $ClO_2^-$

  • Carbonate is $CO_3^{2-}$; a corresponding "carbonite" ion ($CO_2^{2-}$) does not exist in nature.

Nomenclature of Transition Metal Ionic Compounds

• Transition metal cations possess variable oxidation states and can exist with different positive charges (e.g., iron can exist as either $Fe^{2+}$ or $Fe^{3+}$).

• Due to this variability, the specific ionic charge must be explicitly specified in the compound name using Roman numerals in parentheses following the metal name.

• Converting Transition Metal Formulas to Names:

  • Determine the overall negative charge provided by the anion.

  • Calculate the oxidation state (charge) of the transition metal required to achieve an overall neutral charge of $0$.

  • Express that charge as a Roman numeral.

  • Example: In $CoCl_2$, chloride provides a total charge of $2 \times (-1) = -2$. To balance, cobalt must possess a $+2$ charge. The name is cobalt (II) chloride.

  • Example: In $Cu_2O$, oxide provides a charge of $-2$. To balance, the two copper ions must together provide $+2$, giving each copper a $+1$ charge. The name is copper (I) oxide.

• Converting Transition Metal Names to Formulas:

  • The Roman numeral directly states the charge of the cation rather than its group location on the periodic table.

  • Balance this positive charge with the known charge of the anion.

  • Example: Iron (III) oxide denotes $Fe^{3+}$ and $O^{2-}$. Balancing $2(+3) + 3(-2) = 0$ gives the formula $Fe_2O_3$.

Nomenclature of Binary Covalent Compounds

• Binary covalent naming rules apply strictly and exclusively to molecular compounds containing nonmetal elements only (no metals present).

• Numerical prefixes designate the number of atoms of each respective element present in the molecule.

• Standard Numerical Prefixes (Atoms 1 through 10):

  • $1$: mono-

  • $2$: di-

  • $3$: tri-

  • $4$: tetra-

  • $5$: penta-

  • $6$: hexa-

  • $7$: hepta-

  • $8$: octa-

  • $9$: nona-

  • $10$: deca-

• Binary Covalent Naming Rules:

  • Elements are listed in order from left to right as they appear across the Periodic Table.

  • The prefix "mono-" is never applied to the first element if only one atom of that element is present.

  • The second element always uses its appropriate numerical prefix and modifies its root ending to "-ide".

  • Examples:

  • $CO_2$ is named carbon dioxide (the single carbon atom drops "mono-"; two oxygen atoms become "dioxide").

  • $N_2O_3$ is named dinitrogen trioxide (two nitrogen atoms become "dinitrogen"; three oxygen atoms become "trioxide").

Nomenclature of Acids

• The hydrogen cation ($H^+$) is termed the "acid ion".

• All acids release or produce $H^+$ cations when dissolved in aqueous solutions.

• Acids are conceptually treated as ionic compounds wherein $H^+$ functions as the cation paired with an anion.

• Naming Binary / Monoatomic Anion Acids:

  • When the anion is monoatomic, the naming convention follows the pattern:     $\text{hydro} + \text{[anion element root/first syllable]} + \text{ic acid}$

  • Examples:

  • $HCl$ is hydrochloric acid

  • $HF$ is hydrofluoric acid

  • $HBr$ is hydroiodic acid (as specifically identified in systematic acid nomenclature references)

• Naming Oxyacids (Acids Derived from Polyatomic Anions):

  • Polyatomic anions ending in "-ate": The anion ending changes to "-ic acid".

  • $HNO_3$ (from nitrate, $NO_3^-$) is nitric acid

  • $H_2SO_4$ (from sulfate, $SO_4^{2-}$) is sulfuric acid

  • $H_2CO_3$ (from carbonate, $CO_3^{2-}$) is carbonic acid

  • Polyatomic anions ending in "-ite": The anion ending changes to "-ous acid".

  • $H_2SO_3$ (from sulfite, $SO_3^{2-}$) is sulfurous acid

States of Matter and Chemical Sentences

• States of matter are indicated using parenthetical notation positioned immediately after each chemical formula in an equation:

  • $(s)$: solid state

  • $(l)$: liquid state

  • $(g)$: gaseous state

  • $(aq)$: aqueous state (dissolved in water)

• Construction of Chemical Sentences:

  • Standard Sentence Template:     $\text{Reactant 1 } (state) + \text{Reactant 2 } (state) \rightarrow \text{Product 1 } (state) + \text{Product 2 } (state)$

  • Verbal Translation Pattern:     "(state) reactant 1 and (state) reactant 2 produces (state) product 1 and (state) product 2"

  • Example Translation:

  • Chemical Equation:       $Na(s) + H_2O(l) \rightarrow NaOH(aq) + H_2(g)$

  • Spoken/Written Chemical Sentence:       "solid sodium metal and liquid water produces aqueous sodium hydroxide and hydrogen gas"