Empirical and molecular formulas

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The **simplest whole-number ratio** of the atoms of each element in a compound.

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The **actual number** of atoms of each element in one molecule of the compound.

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The molecular formula is a **whole-number multiple** of the empirical formula (molecular = empirical × x).

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**CH_{2}O** — divide every subscript by 6 to get the simplest ratio.

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Treat % as g per 100 g → divide each by A_{r} (n = m/M) → divide by the **smallest** → round / scale to whole numbers.

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**n(C) = n(CO_{2})** — every carbon atom ends up in one CO_{2}.

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**n(H) = 2 × n(H_{2}O)** — each water molecule contains two H atoms.

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By **difference**: subtract the masses of C and H from the sample mass, then divide the leftover by 16.00.

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$x = \dfrac{M_{r}}{\text{empirical formula mass}}$, then multiply every subscript by x.

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Multiply the **whole ratio** by 2 (for .5) or 3 (for .33) to clear it into whole numbers.

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Atoms combine in whole-**number** ratios, which only show up once masses are turned into **moles** (÷ A_{r}).

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An **empirical** formula — ionic compounds have no separate molecules, so the formula is the simplest ratio.