Chapter 4: Chemical Quantities & Aqueous Reactions – Study Notes and Practice Questions

Chapter 4: Chemical Quantities & Aqueous Reactions

Empirical and Molecular Formulas

The empirical formula of a compound represents the simplest whole-number ratio of the elements present. The molecular formula shows the actual number of atoms of each element in a molecule.

• Key Point: To determine the empirical formula, divide the number of moles of each element by the smallest number of moles, then round to the nearest whole number.

• Example: For Hg2(NO3)2, the empirical formula is HgNO3.

Formula:

• $\text{Empirical formula} = \frac{\text{Moles of each element}}{\text{Smallest number of moles}}$

Naming Ionic Compounds

Ionic compounds are named based on the cation and anion present. Transition metals may require Roman numerals to indicate their charge.

• Key Point: The name of the cation comes first, followed by the name of the anion.

• Example: Cu3(PO4)2 is named copper(II) phosphate.

Writing Chemical Formulas

Formulas for compounds are written by balancing the charges of the cation and anion so the overall charge is zero.

• Key Point: Use subscripts to indicate the number of each ion needed to balance the charges.

• Example: For cobalt(III) phosphate: Co3+ and PO43− combine to form CoPO4.

Compound Identification and Naming Practice

Practice naming and writing formulas for common compounds:

• TiO2: titanium(IV) oxide

• SnSO4: tin(II) sulfate

• MnO2: manganese(IV) oxide

• CaCO3: calcium carbonate

• FeCl3: iron(III) chloride

• MgSO4: magnesium sulfate

• Pb(NO3)2: lead(II) nitrate

• PbO2: lead(IV) oxide

Types of Compounds

Compounds can be classified by their composition and bonding:

• Hydrate: Contains water molecules within its crystal structure.

• Double salt: Contains two different salts crystallized together.

• Complex: Contains a central metal atom bonded to ligands.

• Mixed salt: Contains more than one type of cation or anion.

Mole Calculations

The mole is a fundamental unit in chemistry representing $6.022 \times 10^{23}$ entities (Avogadro's number).

• Key Point: To convert between mass and moles, use the molar mass of the compound.

• Formula: $\text{Moles} = \frac{\text{Mass (g)}}{\text{Molar Mass (g/mol)}}$

• Example: To find moles of glucose (C6H12O6) in 31.80 g: $\text{Moles} = \frac{31.80}{180.16}$

Stoichiometry and Mole Relationships

Stoichiometry involves calculating the amounts of reactants and products in chemical reactions.

• Key Point: Use balanced chemical equations to relate moles of different substances.

• Example: Calculate moles of sodium in sodium phosphate (Na3PO4).

Percent Composition

Percent composition is the percentage by mass of each element in a compound.

• Formula: $\%\,\text{Element} = \frac{\text{Mass of element in 1 mol}}{\text{Molar mass of compound}} \times 100$

• Example: Calculate percent composition of sulfur in Al2(SO4)3.

Empirical Formula from Mass Percent

To determine the empirical formula from mass percent:

• Convert percent to grams (assume 100 g sample).

• Convert grams to moles for each element.

• Divide by the smallest number of moles.

• Write the formula using the resulting mole ratios.

Counting Atoms and Molecules

Use Avogadro's number to convert between moles and number of particles.

• Formula: $\text{Number of particles} = \text{Moles} \times 6.022 \times 10^{23}$

• Example: How many molecules of H2C2O4 are in 76.8 g?

Mass Calculations from Molecules

To find the mass from a given number of molecules:

• Convert molecules to moles using Avogadro's number.

• Multiply moles by molar mass to get mass.

• Formula: $\text{Mass (g)} = \text{Moles} \times \text{Molar Mass (g/mol)}$

Bonding and Melting Points

The type of bonding in a compound affects its melting point. Ionic compounds generally have higher melting points than covalent compounds.

• Key Point: Strong ionic bonds result in high melting points; weak covalent bonds result in low melting points.

• Example: NaCl (ionic) has a higher melting point than H2O (covalent).

Classification of Compounds

Compounds can be classified as ionic or molecular (covalent) based on the types of elements present.

• Ionic compounds: Metal + nonmetal (e.g., NaCl, CuO)

• Molecular compounds: Nonmetal + nonmetal (e.g., H2O, NH3)

Table: Classification of Compounds

Compound	Type
CuO	Ionic
Na2O	Ionic
NaNO3	Ionic
SO2	Molecular
NH3	Molecular

Practice Problems

• Determine empirical and molecular formulas from mass percent data.

• Calculate moles, mass, and number of molecules for given compounds.

• Classify compounds as ionic or molecular.

• Predict melting points based on bonding type.

Additional info: These notes expand on worksheet questions by providing definitions, formulas, and examples for key concepts in chemical quantities and compound classification.