Unit Conversions and Significant Figures

Metric and SI Unit Conversions

Unit conversions are essential in chemistry for expressing measurements in appropriate units. Always pay attention to significant figures when performing conversions.

• Length: 1 inch = 2.54 cm (exact)

• Mass: 1 kg = 103 g; 1 g = 103 mg; 1 mg = 103 μg; 1 μg = 103 ng; 1 ng = 103 pg; 1 g = 1015 fg

• Volume: 1 L = 103 mL

Example: Convert 3.51 inches to kilometers.

• First, convert inches to centimeters:

• Then, centimeters to meters:

• Finally, meters to kilometers:

Significant Figures: The number of significant figures in your answer should match the least number in the given data.

Stoichiometry and Mole Calculations

Mole Concept and Avogadro's Number

The mole is a fundamental unit in chemistry representing entities (Avogadro's number).

• Converting grams to moles:

• Converting moles to number of particles:

Example: If you have 9.61 g of (NH4)3PO4, determine the number of oxygen atoms.

• Calculate moles of (NH4)3PO4

• Each formula unit contains 4 oxygen atoms

• Multiply moles by Avogadro's number and by 4

Mass from Number of Molecules

• Formula:

• Convert to required units (e.g., fg, pg, ng) as needed

Chemical Equations and Ionic Equations

Writing and Balancing Equations

Chemical equations must be balanced to obey the law of conservation of mass. Ionic equations show the species that actually participate in the reaction.

• Molecular equation: Shows all reactants and products as compounds

• Total ionic equation: Shows all strong electrolytes as ions

• Net ionic equation: Shows only the species that change during the reaction

Example: Ammonium nitrate + potassium chloride

• Molecular:

• Total ionic:

• Net ionic: No reaction (all ions are spectator ions)

Example: Silver nitrate + iron(III) bromide

• Molecular:

• Total ionic:

• Net ionic:

Percent Composition and Empirical Formulas

Percent Composition

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

• Formula:

Example: A compound contains 2.52 g N and 5.47 g O. Find percent composition.

• Total mass = 2.52 g + 5.47 g = 7.99 g

• %N =

• %O =

Empirical Formula Determination

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

• Convert mass % to grams (assume 100 g sample)

• Convert grams to moles for each element

• Divide by the smallest number of moles to get the ratio

Example: A salt contains 56.58% K, 8.68% C, 34.73% O. Find the empirical formula.

Nomenclature of Ionic and Molecular Compounds

Naming Ionic Compounds

• Cation (metal) + anion (nonmetal)

• Use Roman numerals for transition metals with variable charge

• Polyatomic ions retain their names

Examples:

• Fe3(PO4)2: Iron(II) phosphate

• AgNO3: Silver nitrate

• Au(ClO4)2: Gold(II) perchlorate

• (NH4)2SO4: Ammonium sulfate

• Na2CO3: Sodium carbonate

Naming Acids

• Binary acids: Hydro + root + ic acid (e.g., HCl: hydrochloric acid)

• Oxyacids: -ate to -ic acid, -ite to -ous acid (e.g., HNO3: nitric acid, HNO2: nitrous acid)

Writing Chemical Formulas

• Use the charges of ions to balance the formula

• Examples: Sodium selenide (Na2Se), Ammonium hypobromite ((NH4)BrO), Gold(III) iodate (Au(IO3)3), Iron(II) nitride (Fe3N2), Zinc chloride (ZnCl2)

Solubility of Ionic Compounds

Solubility Rules

Solubility rules help predict whether an ionic compound will dissolve in water.

• Most nitrates (NO3-), ammonium (NH4+), and alkali metal salts are soluble

• Most carbonates (CO32-), phosphates (PO43-), and hydroxides are insoluble except with alkali metals and NH4+

Example: Is FeCO3 soluble? (No, except with alkali metals or NH4+)

Density and Volume Calculations

Density Formula

• Formula:

• To find volume:

Example: A sample has 3.231 μg mass, density 1.92413 g/mL. Find volume in mL.

Isotopes and Atomic Structure

Definition of Isotope

• Isotopes are atoms of the same element with the same number of protons but different numbers of neutrons.

• They have different mass numbers but identical chemical properties.

Classification of Matter

Pure Substances vs. Mixtures

• Pure substance: Has a fixed composition and distinct properties (elements and compounds)

• Mixture: Physical combination of two or more substances; composition can vary

• Example: Sugar (sucrose) is a pure substance

Diatomic Elements

List of Diatomic Elements

• There are seven elements that exist as diatomic molecules in their natural state:

• Hydrogen (H2), Nitrogen (N2), Oxygen (O2), Fluorine (F2), Chlorine (Cl2), Bromine (Br2), Iodine (I2)

Stoichiometry: Limiting Reactant and Yield

Limiting Reactant Problems

In a chemical reaction, the limiting reactant is the substance that is completely consumed first, limiting the amount of product formed.

• Write the balanced chemical equation

• Convert all given masses to moles

• Determine the limiting reactant by comparing mole ratios

• Calculate the amount of product formed

• Find the amount of excess reactant remaining

Example:

• Given masses of N2 and H2, determine which is limiting and calculate mass of NH3 produced

Summary Table: Common Unit Prefixes

Additional info: This study guide covers fundamental concepts in general chemistry, including unit conversions, stoichiometry, nomenclature, empirical formulas, solubility, and basic atomic structure, as inferred from the provided questions.