Acids, Formula Mass, and Composition of Compounds

Chapter 3: Molecules and Compounds

Overview

This chapter introduces the fundamental concepts of acids, the calculation of formula and molar mass, the composition of compounds, and the determination of chemical formulas from experimental data. These topics are essential for understanding the structure and properties of chemical compounds in general chemistry.

Acids

Definition and Properties of Acids

• Acids are molecular compounds that release hydrogen ions (H+) when dissolved in water.

• They are composed of hydrogen (written first in the formula) and one or more nonmetals.

• Example: HCl is a molecular compound that dissociates into H+ (aq) and Cl- (aq) ions in water.

• Acids are characterized by:

  • Sour taste

  • Ability to dissolve certain metals (e.g., Zn, Fe, but not Au, Ag, or Pt)

  • Formulas generally start with H (e.g., HCl, H2SO4)

Types of Acids

• Binary acids: Composed of hydrogen and one nonmetal (contain only two elements).

• Oxyacids: Composed of hydrogen, oxygen, and another element (usually a nonmetal).

Naming Binary Acids

• Write the prefix hydro-.

• Follow with the base name of the nonmetal (root of the nonmetal's name).

• Add the suffix -ic.

• End with the word acid.

Example: HCl (aq) is named hydrochloric acid.

Naming Oxyacids

• The name depends on the ending of the oxyanion (the polyatomic ion containing oxygen):

  • If the oxyanion ends in -ate, change the ending to -ic and add "acid".

  • If the oxyanion ends in -ite, change the ending to -ous and add "acid".

• Do not use the "hydro-" prefix for oxyacids.

Examples:

• HNO3 (aq): Nitrate ion (NO3-) becomes nitric acid.

• HNO2 (aq): Nitrite ion (NO2-) becomes nitrous acid.

Writing Formulas for Acids

• If the name ends in "acid," the formula starts with H.

• For binary acids, use the "hydro-" prefix and combine H with the nonmetal anion.

• For oxyacids, combine H with the appropriate oxyanion to balance the charge.

• The number of H+ ions added must neutralize the charge of the anion.

Example: Sulfuric acid (from sulfate, SO42-): H2SO4

Formula Mass and Molar Mass

Formula Mass

The formula mass (also called molecular mass or molecular weight) is the sum of the atomic masses of all atoms in a chemical formula.

• Calculated as: $\text{Formula mass} = \sum (\text{Number of atoms of element} \times \text{Atomic mass of element})$

• Units: atomic mass units (amu)

Example: For CO2:

• 1 atom C: 12.01 amu

• 2 atoms O: 2 × 16.00 amu = 32.00 amu

• Total: 12.01 + 32.00 = 44.01 amu

Molar Mass

• The molar mass of a compound is numerically equal to its formula mass, but expressed in grams per mole (g/mol).

• It represents the mass of one mole of molecules or formula units.

Example: The molar mass of CO2 is 44.01 g/mol.

Composition of Compounds

Mass Percent Composition

The mass percent composition of an element in a compound is the percentage by mass of that element in the compound.

• Calculated as: Mass % of element X = (Mass of X in 1 mol of compound)$\div$(Mass of 1 mol of compound) $\times$100

• Mass percentages may not always total exactly 100% due to rounding.

Example: In CCl2F2 (chlorofluorocarbon):

• Mass % Cl = $\frac{2 \times \text{molar mass of Cl}}{\text{molar mass of CCl}_2\text{F}_2} \times 100$

• Suppose this gives 58.64% Cl.

Using Mass Percent as a Conversion Factor

• Mass percent can be used to convert between the mass of an element and the mass of a compound.

• Set up as a ratio: $\frac{\text{mass of element}}{100\,\text{g compound}}$

• Example: If CCl2F2 is 58.64% Cl, then 100 g of CCl2F2 contains 58.64 g Cl.

Determining a Chemical Formula from Experimental Data

Empirical Formula

The empirical formula is the simplest whole-number ratio of atoms of each element in a compound. It can be determined from elemental analysis or mass percent composition.

• Steps to determine empirical formula:

  1. Convert percentages to grams (assume 100 g sample if only percentages are given).

  2. Convert grams to moles using molar mass.

  3. Write a pseudoformula using the calculated moles as subscripts.

  4. Divide all subscripts by the smallest number of moles to get the simplest ratio.

  5. If necessary, multiply all subscripts by a small whole number to obtain whole numbers (e.g., if a subscript is 0.5, multiply all by 2).

Example: A compound contains 24.5 g N and 70.0 g O. Find the empirical formula.

1. Convert to moles: $\text{moles N} = \frac{24.5}{14.01}$, $\text{moles O} = \frac{70.0}{16.00}$

2. Write pseudoformula: NxOy

3. Divide by smallest number of moles to get whole-number ratio.

Molecular Formula

The molecular formula is a whole-number multiple of the empirical formula. To determine it, you need the empirical formula and the molar mass of the compound.

• Relationship: $\text{Molecular formula} = (\text{Empirical formula})_n$

• Where $n = \frac{\text{Molar mass}}{\text{Empirical formula molar mass}}$

Example: Fructose has empirical formula CH2O and molar mass 180.2 g/mol.

• Empirical formula mass: 12.01 + 2(1.01) + 16.00 = 30.03 g/mol

• $n = \frac{180.2}{30.03} = 6$

• Molecular formula: C6H12O6

Summary Table: Types of Acids

Type of Acid	Composition	Naming Pattern	Example
Binary Acid	H + nonmetal	hydro- + base name + -ic acid	HCl (hydrochloric acid)
Oxyacid	H + polyatomic oxyanion	If -ate: base name + -ic acid If -ite: base name + -ous acid	HNO3 (nitric acid), HNO2 (nitrous acid)

Key Equations

• Formula mass: $\sum (\text{Number of atoms} \times \text{Atomic mass})$

• Mass percent: $\frac{\text{Mass of element in 1 mol}}{\text{Molar mass of compound}} \times 100$

• Molecular formula: $(\text{Empirical formula})_n$, $n = \frac{\text{Molar mass}}{\text{Empirical formula mass}}$