Introduction to Chemistry: Study Guide for Chapters 5–7 (Atoms, Compounds, Chemical Reactions, and Stoichiometry)

Atoms, Elements, and Compounds

Types of Atoms in Chemical Formulas

Understanding the composition of chemical formulas is fundamental in chemistry. Each formula represents the types and numbers of atoms present in a compound.

• Atoms: The smallest unit of an element, retaining its chemical properties.

• Example: In the formula (NH4)2HPO4, count each type of atom:

  • Nitrogen (N): 2 atoms

  • Hydrogen (H): 9 atoms (8 from NH4 and 1 from HPO4)

  • Phosphorus (P): 1 atom

  • Oxygen (O): 4 atoms

Writing and Interpreting Chemical Formulas

Chemical formulas indicate the elements and their ratios in a compound.

• Example: A compound with ten oxygen atoms and four phosphorus atoms: P4O10

• Ionic Compounds: Formed from metals and nonmetals. Example: Aluminum and oxygen form Al2O3; potassium and oxygen form K2O.

Naming Compounds

Compounds are named according to specific rules based on their composition.

• CaCO3: Calcium carbonate

• Al2(SO4)3: Aluminum sulfate

• KOH: Potassium hydroxide

Matching Names and Formulas

It is important to correctly match compound names with their formulas.

• Copper(I) nitrate: CuNO3

• Copper(II) nitrate: Cu(NO3)2

• Copper(II) nitrite: Cu(NO2)2

Naming Molecular Compounds

Molecular compounds are named using prefixes to indicate the number of atoms.

• SO3: Sulfur trioxide

• SF5: Sulfur pentafluoride

• Heptaphosphorus octafluoride: P7F8

Classification of Substances

Types of Elements and Compounds

Substances can be classified as atomic elements, molecular elements, or molecular compounds.

• Atomic element: Exists as single atoms (e.g., Krypton).

• Molecular element: Exists as molecules of the same element (e.g., O2).

• Molecular compound: Molecules composed of different elements (e.g., SO2).

Mass, Moles, and Stoichiometry

Counting Atoms Using Mass and Moles

The mole is a fundamental unit for counting atoms and molecules in chemistry.

• Mole: $1~\text{mol} = 6.022 \times 10^{23}$ entities (Avogadro's number).

• Example: One mole of nitrogen gas contains $2 \times 6.022 \times 10^{23}$ nitrogen atoms (since N2).

Empirical Formulas

The empirical formula shows the simplest whole-number ratio of atoms in a compound.

• Example: C2H3O2 could be an empirical formula.

Calculating Number of Atoms and Moles

• Number of atoms: $\text{Number of atoms} = \text{moles} \times 6.022 \times 10^{23}$

• Example: Number of mercury atoms in 5.8 mol: $5.8 \times 6.022 \times 10^{23}$

Calculating Molar Mass

Molar mass is the mass of one mole of a substance, expressed in grams per mole.

• Potassium nitrate (KNO3): $M = 39.10 + 14.01 + (3 \times 16.00) = 101.11~\text{g/mol}$

• Diboron tetrachloride (B2Cl4): $M = (2 \times 10.81) + (4 \times 35.45) = 163.43~\text{g/mol}$

• Calcium nitrate (Ca(NO3)2): $M = 40.08 + 2 \times [14.01 + (3 \times 16.00)] = 164.10~\text{g/mol}$

Mass Percent and Empirical Formula Calculations

• Mass percent: $\text{Mass percent} = \frac{\text{mass of element}}{\text{total mass}} \times 100\%$

• Empirical formula: Use masses of elements to determine the simplest ratio.

• Example: If decomposition produces 1.78 g N and 4.05 g O, calculate moles and find the ratio.

Chemical Reactions and Equations

Evidence of Chemical Reactions

Chemical reactions can be identified by observable changes.

• Formation of a new solid or gas

• Color change

• Emission of light

• Gas formation

Balancing Chemical Equations

Balanced equations have equal numbers of each atom on both sides.

• Example: $\text{AgNO}_3 + 2\text{NH}_4\text{Cl} \rightarrow \text{AgCl} + (\text{NH}_4)_2\text{NO}_3$

• Example: $\text{HNO}_3 + \text{NaHSO}_3 \rightarrow \text{NaNO}_3 + \text{H}_2\text{O} + \text{SO}_2$

Types of Chemical Reactions

• Double displacement: $\text{AB} + \text{CD} \rightarrow \text{AD} + \text{CB}$

• Single displacement: $\text{A} + \text{BC} \rightarrow \text{AC} + \text{B}$

Electrolytes and Solubility

Electrolytes are substances that dissociate into ions in solution.

• Strong electrolytes: Completely dissociate in water.

• Solubility: Soluble compounds dissolve completely; insoluble compounds do not.

Precipitation Reactions

Precipitation occurs when an insoluble solid forms from mixing solutions.

• Example: $\text{Pb(NO}_3)_2~(aq) + 2\text{KCl}~(aq) \rightarrow \text{PbCl}_2~(s) + 2\text{KNO}_3~(aq)$

• Spectator ions: Ions that do not participate in the reaction (e.g., K+, NO3-).

Complete Ionic and Net Ionic Equations

Complete ionic equations show all ions present; net ionic equations show only those involved in the reaction.

• Example: In the reaction above, Pb2+ and Cl- form the precipitate; K+ and NO3- are spectator ions.

Identifying Double Displacement Reactions

Double displacement reactions involve the exchange of ions between two compounds.

• Example:

  • KCl(aq) + AgNO3(aq) → AgCl(s) + KNO3(aq)

  • Na2SO4(aq) + BaCl2(aq) → BaSO4(s) + 2NaCl(aq)

  • H2SO4(aq) + 2NaOH(aq) → Na2SO4(aq) + 2H2O(l)

Tables

Classification Table: Types of Substances

Substance	Classification
Krypton	Atomic element
SO2	Molecular compound
O2	Molecular element

Solubility Table (Sample)

Compound	Solubility in Water
Na2SO4	Soluble
PbCl2	Insoluble (forms precipitate)

Common Types of Chemical Reactions

Type	General Equation	Example
Double Displacement	AB + CD → AD + CB	KCl + AgNO3 → AgCl + KNO3
Single Displacement	A + BC → AC + B	Mg + 2HCl → MgCl2 + H2

Additional info: Some explanations and tables have been expanded for clarity and completeness, based on standard introductory chemistry curriculum.