Question

Which are correct formulas for ionic compounds? For those that are not correct, give the correct formula and justify your answer. $$ \begin{array}{ll}{\text { a. AlCl }} & {\text { c. BaOH }_{2}} \\ {\text { b. } \mathrm{Na}_{3} \mathrm{SO}_{4}} & {\text { d. Fe_ } \mathrm{O}}\end{array} $$

Short answer

The correct formulas for the given ionic compounds are: a. AlCl is incorrect. The correct formula is \( \mathrm{AlCl}_{3} \). b. Na3SO4 is correct. c. BaOH2 is correct. d. FeO is incorrect without specifying the oxidation state of Fe. The possible correct formulas are \( \mathrm{Fe}_{2} \mathrm{O}_{3} \) or \( \mathrm{FeO} \), depending on the charge of the iron ion.

Step-by-step solution

Identify the ions involved

First, let's identify the ions involved in each of the given formulas. a. AlCl - Aluminum ion (Al^3+) and Chloride ion (Cl^-) b. Na3SO4 - Sodium ion (Na^+) and Sulfate ion (SO4^2-) c. BaOH2 - Barium ion (Ba^2+) and Hydroxide ion (OH^-) d. FeO - Iron ion (Fe) and Oxide ion (O^2-)

Check if the charges are balanced

Now, let's examine the given chemical formulas and see if the charges of the ions are balanced. a. AlCl - Al^3+ and Cl^- (not balanced, correct formula \( \: \mathrm{AlCl}_{3} \)) b. Na3SO4 - Na^+ and SO4^2- (balanced) c. BaOH2 - Ba^2+ and OH^- (balanced) d. FeO - Fe and O^2- (missing charge on Fe, correct formula \( \: \mathrm{Fe}_{2} \mathrm{O}_{3} \) or \( \: \mathrm{FeO} \) depending on the oxidation state of Fe)

Check if the subscripts are simplified

In this step, we will verify if the subscripts used in the formulas are simplified. a. AlCl3 - The subscripts are simplified. b. Na3SO4 - The subscripts are simplified. c. BaOH2 - The subscripts are simplified. d. Fe2O3 or FeO - The subscripts are simplified in both cases.

Final answers

Now, let's summarize our findings. a. AlCl is incorrect. The correct formula is \( \mathrm{AlCl}_{3} \). b. Na3SO4 is correct. c. BaOH2 is correct. d. FeO is incorrect without specifying the oxidation state of Fe. The possible correct formulas are \( \mathrm{Fe}_{2} \mathrm{O}_{3} \) or \( \mathrm{FeO} \), depending on the charge of the iron ion.

Key concepts

Chemical Formulas

A chemical formula is a way of representing the elements in a compound. It shows the types and numbers of atoms in the simplest ratio of the compound. In ionic compounds, the formula reflects the balance of charges between the cations (positively charged ions) and anions (negatively charged ions). For example, in aluminum chloride, AlCl, the correct formula is actually \( \mathrm{AlCl}_{3} \). This is because three chloride ions, each with a \(-1\) charge, are needed to balance the \(+3\) charge of one aluminum ion.

Writing correct chemical formulas is essential because it tells us the precise proportions of each ion in the compound. Incorrect formulas can lead to misunderstanding about the nature of the compound. Therefore, understanding how to create and interpret these formulas is a fundamental chemistry skill.

If you see a formula that doesn't seem to uphold a balance of charges, it is likely incorrect and needs further examination.

Ion Charges

Ion charges are fundamental to understanding chemical formulas for ionic compounds. An ion is an atom or molecule that has gained or lost one or more electrons, thereby acquiring a charge. The charge on an ion is crucial because it determines how ions combine to form neutral compounds.

The formula of an ionic compound must reflect a balance of these charges. In our examples, the aluminum ion is \( \mathrm{Al}^{3+} \) and the chloride ion is \( \mathrm{Cl}^{-} \). To balance the charges, three chloride ions are needed for each aluminum ion, resulting in \( \mathrm{AlCl}_{3} \).

Similarly, for sodium sulfate, the sodium ion is \( \mathrm{Na}^{+} \), and the sulfate ion is \( \mathrm{SO}_{4}^{2-} \). Two sodium ions are needed to balance the charge of one sulfate ion, hence the correct formula \( \mathrm{Na}_{2}\mathrm{SO}_{4} \). Charge balance is key to determining the correct formula of an ionic compound.

Oxidation States

Oxidation states offer a way to keep track of electrons in compounds and are especially helpful for understanding transition metals. Each element in a compound is assigned an oxidation state, which helps us to predict the possible ionic forms an element may take.

In the case of iron oxide, iron can exist in multiple oxidation states, such as \(+2\) or \(+3\). These different states lead to different compounds, like \( \mathrm{FeO} \) or \( \mathrm{Fe}_{2}\mathrm{O}_{3} \), each reflecting a different relative amount of iron to oxygen.

Understanding oxidation states allows scientists and chemists to write correct formulas for ionic compounds by using the known charges of ions. This ensures the charges are balanced, leading to a neutrally charged compound. Clarifying oxidation states is particularly important for dealing with elements like transition metals, which can have multiple stable oxidation states.

Subscript Simplification

Subscript simplification in chemical formulas involves ensuring that the ion ratios are expressed in the simplest form. In a balanced chemical formula, subscripts should reflect the smallest whole number ratio of ions that result in charge neutrality.

For example, in the formula \( \mathrm{AlCl}_{3} \), the subscripts "3" indicates that three chloride ions are needed per aluminum ion. There is no need to simplify this further as the ratio directly balances the \(+3\) charge from aluminum with three \(-1\) charges from each chloride ion.

In another example, \( \mathrm{Ba(OH)}_{2} \), the ratio of \( \mathrm{Ba}^{2+} \) to \( \mathrm{OH}^{-} \) is naturally balanced by having two hydroxide ions for each barium ion, giving a neutral compound. Simplification is not necessary since the ratio of ions already represents the simplest whole numbers required for charge balance.