A chemical formula serves as a universal language in science, representing the exact number and type of atoms that make up a compound. For acids, this formula specifies the components of a substance that, when dissolved in water, yields hydrogen ions (\(\text{H}^+\)). This release of protons is the defining characteristic that gives acids their unique chemical properties. Understanding how to write these formulas is a fundamental skill that connects a compound’s name directly to its molecular structure. This guide provides a systematic method for converting the common names of acids into their correct chemical formulas.
Understanding Acid Classification
Acids are broadly categorized into two groups based on their elemental composition, a distinction that dictates the formula writing process. The first group consists of binary acids, which are compounds formed solely from hydrogen and one other non-metallic element. Examples include hydrochloric acid (\(\text{HCl}\)) and hydrobromic acid (\(\text{HBr}\)). These acids do not contain oxygen atoms in their structure.
The second and larger group is known as oxyacids, which always contain oxygen in addition to hydrogen and a third non-metallic element. These acids are characterized by the presence of a polyatomic ion, a charged group of atoms acting as a single unit. Nitric acid (\(\text{HNO}_3\)) and sulfuric acid (\(\text{H}_2\text{SO}_4\)) are common examples of oxyacids.
Determining Formulas for Binary Acids
The systematic name for a binary acid provides all the necessary information to determine its chemical formula. These names consistently begin with the prefix “hydro-” and end with the suffix “-ic acid.” The root of the name points directly to the non-metal element that combines with hydrogen; for instance, the “fluor-” root in hydrofluoric acid indicates the presence of the fluoride ion.
The first step involves identifying the specific anion that forms the acid. The non-metal is always in its ionic form, carrying a specific negative charge that corresponds to its position on the periodic table. The fluoride ion, for example, has a charge of 1- (\(\text{F}^-\)).
The resulting formula must be electrically neutral overall, meaning the total positive charge must equal the total negative charge. Since each hydrogen atom contributes a single positive charge (\(\text{H}^+\)), the number of hydrogen atoms required is equal to the magnitude of the anion’s negative charge. For hydrofluoric acid, one \(\text{H}^+\) is needed to balance the one \(\text{F}^-\), resulting in the formula \(\text{HF}\). Similarly, hydrosulfuric acid requires two \(\text{H}^+\) ions to balance the 2- charge of the sulfide ion (\(\text{S}^{2-}\)), leading to the formula \(\text{H}_2\text{S}\).
Determining Formulas for Oxyacids
Writing the formula for an oxyacid requires the intermediate step of identifying the correct polyatomic ion. Unlike binary acids, oxyacid names do not use the “hydro-” prefix, and the ending suffix relates directly to the composition of the oxygen-containing polyatomic ion.
The key to this process is a simple, inverse relationship between the acid’s suffix and the polyatomic ion’s suffix. If the acid name ends in “-ic acid,” the polyatomic ion must end in “-ate.” For example, sulfuric acid is derived from the sulfate ion, and nitric acid is derived from the nitrate ion.
Conversely, if the acid name ends with the suffix “-ous acid,” the polyatomic ion must end in “-ite.” Sulfurous acid is formed from the sulfite ion, and nitrous acid from the nitrite ion. The “-ous” form always indicates one less oxygen atom than its corresponding “-ic” or “-ate” counterpart, while the charge remains the same.
Once the correct polyatomic ion is identified, its chemical formula and charge must be known or looked up. For instance, the sulfate ion is \(\text{SO}_4^{2-}\), and the sulfite ion is \(\text{SO}_3^{2-}\). This polyatomic ion then acts as the single anion that must be neutralized by the hydrogen ions.
The principle of electrical neutrality applies here. The total positive charge from the hydrogen ions must exactly cancel the total negative charge of the polyatomic ion. The sulfate ion (\(\text{SO}_4^{2-}\)), with its 2- charge, requires two \(\text{H}^+\) ions to achieve neutrality, resulting in the formula \(\text{H}_2\text{SO}_4\).
The nitrate ion (\(\text{NO}_3^-\)), carrying a 1- charge, only needs one \(\text{H}^+\) ion, yielding \(\text{HNO}_3\). The final formula is constructed by writing the hydrogen atoms first, followed by the entire polyatomic ion. Subscripts must reflect the number of atoms or ions needed for a net zero charge, allowing for the accurate conversion of names like perchloric acid (from perchlorate, \(\text{ClO}_4^-\)) to \(\text{HClO}_4\).
Key Examples and Common Acid Formulas
Applying the rules of nomenclature allows for the accurate determination of formulas for several common acids. Sulfuric acid, derived from the sulfate ion (\(\text{SO}_4^{2-}\)), requires two hydrogen atoms to balance its charge, resulting in the formula \(\text{H}_2\text{SO}_4\). Another important oxyacid is phosphoric acid, which comes from the phosphate ion (\(\text{PO}_4^{3-}\)). Because the phosphate ion carries a 3- charge, three \(\text{H}^+\) ions are necessary, establishing the formula as \(\text{H}_3\text{PO}_4\).
Hydrochloric acid, a common binary acid, is formed by combining the single \(\text{H}^+\) with the single \(\text{Cl}^-\) ion, leading simply to \(\text{HCl}\). One notable acid that slightly deviates from the standard pattern is acetic acid (\(\text{HC}_2\text{H}_3\text{O}_2\)), the main component of vinegar. It is an organic acid, meaning it is carbon-based, and only the hydrogen atom written first is acidic. Despite its complex structure, the charge-balancing principles still apply to the overall acetate ion.