Percentage Composition — Revision Notes

⚡ 30-Second Revision

Definition — Mass percentage of each element in a compound.
Formula — $\% \text{Element} = \frac{\text{Mass of element in 1 mole}}{\text{Molar mass of compound}} \times 100$
Basis — Law of Definite Proportions.
Empirical Formula — Simplest whole-number ratio of atoms.

- Steps: % to mass $\rightarrow$ mass to moles $\rightarrow$ divide by smallest mole $\rightarrow$ whole numbers.

Molecular Formula — Actual number of atoms.

- $\text{Molecular Formula} = (\text{Empirical Formula})_n$, where $n = \frac{\text{Molar Mass}}{\text{Empirical Formula Mass}}$.

Key — Accurate molar mass calculation.

2-Minute Revision

Percentage composition is a fundamental concept in chemistry that tells us the mass contribution of each element in a compound, expressed as a percentage. It's based on the Law of Definite Proportions, meaning a pure compound always has the same elemental mass ratios.

To calculate it, you find the molar mass of the compound, then the total mass of a specific element within one mole of that compound, and finally divide the element's mass by the compound's molar mass, multiplying by 100.

For example, in $H_2O$, oxygen is about 88.8% by mass. This concept is crucial for determining the empirical formula (simplest whole-number ratio of atoms) from experimental data. You convert mass percentages to moles, find the simplest mole ratio, and then convert to whole numbers.

If the molar mass is also known, the molecular formula (actual number of atoms) can be found by determining how many empirical formula units make up the molecule. Always ensure accurate molar mass calculations and careful handling of ratios.

5-Minute Revision

Percentage composition is the mass percentage of each element in a chemical compound. It's a direct application of the Law of Definite Proportions, which states that a pure compound always contains elements in fixed mass ratios. For example, if you have a compound like $C_2H_5OH$ (ethanol), you'd calculate its molar mass: $(2 \times 12) + (6 \times 1) + (1 \times 16) = 24 + 6 + 16 = 46,\text{g/mol}$.

Then, for each element:

Carbon — Mass of C = $2 \times 12 = 24,\text{g}$. %C = $(24/46) \times 100 = 52.17$
Hydrogen — Mass of H = $6 \times 1 = 6,\text{g}$. %H = $(6/46) \times 100 = 13.04$
Oxygen — Mass of O = $1 \times 16 = 16,\text{g}$. %O = $(16/46) \times 100 = 34.78$

(Sum $\approx 100$).

This concept is vital for determining empirical and molecular formulas. If you're given percentage composition, assume a 100g sample to convert percentages to masses. Then, convert these masses to moles by dividing by atomic masses.

Divide all mole values by the smallest mole value to get a preliminary ratio. If not whole numbers, multiply by the smallest integer to get whole numbers for the empirical formula. If the molar mass is also given, calculate the empirical formula mass (EFM).

The ratio $n = \text{Molar Mass} / \text{EFM}$ gives the factor by which to multiply the empirical formula subscripts to get the molecular formula. For instance, if empirical formula is $CH_2$ and molar mass is $42,\text{g/mol}$, EFM = 14 g/mol.

$n = 42/14 = 3$. Molecular formula = $(CH_2)_3 = C_3H_6$. Remember to be precise with atomic masses and calculations.

Prelims Revision Notes

Percentage composition is the mass percentage of each element in a compound. It's a key quantitative aspect of chemical compounds.

1. Calculation from Chemical Formula:

* **Step 1: Find Molar Mass of Compound ($M_{compound}$)**: Sum of atomic masses of all atoms in the formula. E.g., for $CO_2$, $M_{CO_2} = 12 + (2 \times 16) = 44,\text{g/mol}$. * **Step 2: Find Total Mass of Element in One Mole ($M_{element}$)**: Multiply atomic mass of the element by its subscript in the formula.

E.g., for O in $CO_2$, $M_O = 2 \times 16 = 32,\text{g}$. * Step 3: Calculate Percentage: $\% \text{Element} = \frac{M_{element}}{M_{compound}} \times 100$. E.g., %O in $CO_2 = (32/44) \times 100 = 72.

73\%$.

2. Law of Definite Proportions:

* Underpins percentage composition. States that a pure compound always contains elements in fixed mass ratios, regardless of source.

3. Determining Empirical Formula from Percentage Composition:

* Step 1: Assume 100g sample: Converts % directly to mass (e.g., 60% C $\rightarrow$ 60g C). * Step 2: Convert Mass to Moles: Divide mass of each element by its atomic mass. * Step 3: Find Smallest Mole Ratio: Divide all mole values by the smallest mole value obtained. * Step 4: Convert to Whole Numbers: If ratios are not whole numbers, multiply all by the smallest integer to get whole numbers. These are the subscripts for the empirical formula.

4. Determining Molecular Formula from Empirical Formula and Molar Mass:

* Step 1: Calculate Empirical Formula Mass (EFM): Sum of atomic masses in the empirical formula. * Step 2: Find 'n': $n = \frac{\text{Molar Mass of Compound}}{\text{EFM}}$ (n must be a whole number). * Step 3: Molecular Formula: Multiply all subscripts in the empirical formula by 'n'. Molecular Formula = $(EF)_n$.

5. Hydrated Salts:

* Treat water of crystallization ($xH_2O$) as a component. Calculate its total mass and use the percentage formula. E.g., for $CuSO_4 \cdot 5H_2O$, % water = $(5 \times M_{H_2O} / M_{CuSO_4 \cdot 5H_2O}) \times 100$.

Common Mistakes to Avoid:

* Incorrect molar mass calculation. * Confusing mass percentage with mole percentage. * Rounding errors in intermediate steps for empirical formula determination.

Vyyuha Quick Recall

For 'Percentage Composition' and 'Empirical Formula':

Percent Mass To Moles, Divide Smallest, Multiply Whole.

Percent: Start with percentage composition.
Mass: Convert percentages to mass (assume 100g).
To Moles: Convert mass to moles (divide by atomic mass).
Divide Smallest: Divide all mole values by the smallest mole value.
Multiply Whole: Multiply by a factor to get whole numbers (if necessary) for the empirical formula subscripts.