General Chemistry: Naming, Formulas, and Mass Calculations Study Guide

Polyatomic Ions and Their Memorization

Importance of Polyatomic Ions

Polyatomic ions are ions composed of two or more atoms covalently bonded, carrying a net charge. Memorizing their names, formulas, and charges is essential for writing correct chemical formulas and equations.

• Common Polyatomic Ions: Examples include nitrate ($NO_3^-$), sulfate ($SO_4^{2-}$), ammonium ($NH_4^+$), and carbonate ($CO_3^{2-}$).

• Application: Used in naming ionic compounds and predicting products in chemical reactions.

Prefixes for Molecular Compounds and Hydrates

Systematic Naming Using Prefixes

Prefixes indicate the number of atoms in molecular compounds and the number of water molecules in hydrates.

• Common Prefixes: mono- (1), di- (2), tri- (3), tetra- (4), penta- (5), hexa- (6), hepta- (7), octa- (8), nona- (9), deca- (10)

• Example: Carbon dioxide ($CO_2$), dinitrogen tetroxide ($N_2O_4$), copper(II) sulfate pentahydrate ($CuSO_4 \, 5H_2O$)

Ionic vs. Covalent Bonds

Differences Between Bond Types

Ionic and covalent bonds are two primary types of chemical bonds, differing in electron transfer and sharing.

• Ionic Bonds: Formed by the transfer of electrons from a metal to a nonmetal, resulting in oppositely charged ions (e.g., $NaCl$).

• Covalent Bonds: Formed by the sharing of electrons between two nonmetals (e.g., $H_2O$).

Empirical, Molecular, and Structural Formulas

Types of Chemical Formulas

Chemical formulas represent the composition of compounds in different ways.

• Empirical Formula: Shows the simplest whole-number ratio of atoms (e.g., $CH_2O$ for glucose).

• Molecular Formula: Shows the actual number of each atom in a molecule (e.g., $C_6H_{12}O_6$ for glucose).

• Structural Formula: Depicts the arrangement of atoms and bonds (e.g., Lewis structures).

Classification of Pure Substances

Types of Pure Substances

Pure substances can be classified as elements or compounds.

• Element: A substance made of only one type of atom (e.g., $O_2$, $Fe$).

• Compound: A substance composed of two or more different elements chemically bonded (e.g., $H_2O$, $NaCl$).

Writing Formulas for Ionic Compounds

Steps for Writing Ionic Formulas

Formulas for ionic compounds are written by balancing the charges of the cations and anions.

• Write the symbol and charge for the cation and anion.

• Balance the total positive and negative charges to zero.

• Use subscripts to indicate the number of each ion needed.

• Example: $Al^{3+}$ and $O^{2-}$ combine to form $Al_2O_3$.

Naming Ionic Compounds

Rules for Naming

Ionic compounds are named by stating the cation first, followed by the anion.

• For metals with variable charges, indicate the charge with Roman numerals (e.g., iron(III) chloride for $FeCl_3$).

• Polyatomic ions retain their names (e.g., sodium sulfate for $Na_2SO_4$).

Naming Hydrates

Hydrate Nomenclature

Hydrates are ionic compounds with water molecules incorporated into their structure.

• Name the ionic compound, then add a prefix and the word "hydrate" to indicate the number of water molecules.

• Example: $CuSO_4 \, 5H_2O$ is copper(II) sulfate pentahydrate.

Naming Molecular Compounds

Rules for Molecular Compound Names

Molecular compounds (nonmetal + nonmetal) use prefixes to indicate the number of each atom.

• The first element keeps its name; the second element ends in "-ide".

• Prefixes are used for both elements, except "mono-" is often omitted for the first element.

• Example: $N_2O_5$ is dinitrogen pentoxide.

Naming Acids

Acid Nomenclature

Acids are named based on the anion they contain.

• If the anion ends in "-ide": use "hydro-...-ic acid" (e.g., $HCl$ is hydrochloric acid).

• If the anion ends in "-ate": use "...-ic acid" (e.g., $H_2SO_4$ is sulfuric acid).

• If the anion ends in "-ite": use "...-ous acid" (e.g., $H_2SO_3$ is sulfurous acid).

Calculating Formula Mass and Molar Mass

Definitions and Calculations

Formula mass and molar mass are measures of the mass of a compound's formula unit or mole, respectively.

• Formula Mass: Sum of atomic masses of all atoms in a formula unit (in amu).

• Molar Mass: Mass of one mole of a substance (in g/mol).

• Calculation: Add the atomic masses of each element multiplied by the number of atoms present.

• Example: For $H_2O$, $2(1.01) + 16.00 = 18.02$ g/mol.

Calculating Mass Percent and Its Use as a Conversion Factor

Mass Percent Calculation

Mass percent expresses the mass of an element in a compound as a percentage of the total mass.

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

• Application: Used to determine how much of an element is present in a given mass of compound.

• Example: In $H_2O$, mass percent of H is $\frac{2.02}{18.02} \times 100\% = 11.2\%$.

Using Conversion Factors from Chemical Formulas

Stoichiometric Relationships

Chemical formulas provide conversion factors between moles of compounds and moles of constituent elements.

• Example: $1$mol$H_2O$contains $2$ mol H and $1$ mol O.

• These relationships are used in stoichiometry to convert between amounts of substances.

Calculating Grams of an Element in a Given Mass of Compound

Stepwise Calculation

To find the mass of an element in a sample of a compound, use the mass percent or mole ratios.

• Calculate moles of compound from given mass.

• Use the chemical formula to find moles of the element.

• Convert moles of the element to grams using its molar mass.

• Example: How many grams of H are in 36.04 g of $H_2O$?

  1. Find moles of $H_2O$: $\frac{36.04\,g}{18.02\,g/mol} = 2.00$ mol

  2. Moles of H: $2.00$ mol $H_2O \times \frac{2\,mol\,H}{1\,mol\,H_2O} = 4.00$ mol H

  3. Grams of H: $4.00$ mol $\times 1.01\,g/mol = 4.04$ g