Chapter 1: Introduction – Matter and Measurement

Study of Chemistry

Chemistry is the scientific study of matter, its properties, composition, and the changes it undergoes. Understanding the classification and measurement of matter is foundational for all chemical studies.

• Definition of Chemistry: The science that investigates the composition, structure, properties, and changes of matter.

• States of Matter: Matter exists in three primary states: solid (definite shape and volume), liquid (definite volume, variable shape), and gas (variable shape and volume).

• Physical vs. Chemical Properties:

  • Physical properties can be observed without changing the substance's identity (e.g., melting point, density).

  • Chemical properties describe a substance's ability to undergo chemical changes (e.g., flammability, reactivity).

• Classification of Matter:

  • Element: Pure substance made of one type of atom (e.g., O2).

  • Compound: Pure substance composed of two or more elements chemically combined (e.g., H2O).

  • Homogeneous mixture: Uniform composition throughout (e.g., saltwater).

  • Heterogeneous mixture: Non-uniform composition (e.g., salad, sand in water).

• Intensive vs. Extensive Properties:

  • Intensive properties do not depend on the amount of substance (e.g., density, boiling point).

  • Extensive properties depend on the amount of substance (e.g., mass, volume).

Measurement and Units

• SI Base Units: The International System of Units (SI) uses standard units for measurement:

  • Length: meter (m)

  • Mass: kilogram (kg)

  • Time: second (s)

  • Temperature: kelvin (K)

  • Amount of substance: mole (mol)

• Metric Prefixes: Used to express multiples or fractions of units (e.g., kilo-, centi-, milli-).

• Unit Conversions: Ability to convert between metric units (e.g., cm to m, g to kg).

• Temperature Conversions:

  • Celsius to Kelvin:

  • Kelvin to Celsius:

Density

Density is a physical property defined as mass per unit volume.

• Formula:

• Applications: Used to identify substances and solve for mass or volume when density is known.

• Example: If a sample has a mass of 10.0 g and a volume of 2.0 mL, its density is .

Significant Figures and Scientific Notation

• Significant Figures (Sig Figs): Digits in a measurement that are known with certainty plus one estimated digit.

  • Nonzero digits are always significant.

  • Captive zeros (between nonzero digits) are significant.

  • Leading zeros (before nonzero digits) are not significant.

  • Trailing zeros are significant only if a decimal point is present.

• Scientific Notation: Expresses numbers as a coefficient times a power of ten (e.g., ).

• Determining Significant Figures: Count according to the rules above.

• Calculations: Use the correct number of significant figures in answers:

  • Multiplication/Division: Result has as many sig figs as the measurement with the fewest sig figs.

  • Addition/Subtraction: Result has as many decimal places as the measurement with the fewest decimal places.

Dimensional Analysis

Dimensional analysis is a method for converting between units using conversion factors.

• Conversion Factors: Ratios that express how many of one unit are equal to another unit (e.g., ).

• Process: Multiply by conversion factors so that units cancel, leaving the desired unit.

• Multistep Problems: Chain multiple conversion factors as needed.

• Example: Convert 25.0 inches to centimeters:

Chapter 2: Atoms, Molecules, and Ions

Atomic Theory & Discovery of Atomic Structure

The modern understanding of atomic structure is based on key experiments and discoveries by several scientists.

• J. J. Thomson: Discovered the electron using the cathode ray tube experiment.

• Millikan: Measured the charge of the electron with the oil drop experiment.

• Rutherford: Discovered the atomic nucleus via the gold foil experiment.

• Chadwick: Discovered the neutron.

Structure of the Atom

• Atomic Number (Z): Number of protons in the nucleus; defines the element.

• Mass Number (A): Total number of protons and neutrons in the nucleus.

• Determining Subatomic Particles:

  • Protons = Atomic number (Z)

  • Neutrons = Mass number (A) – Atomic number (Z)

  • Electrons = Protons (for neutral atoms)

• Isotopes: Atoms of the same element with different numbers of neutrons (different mass numbers).

• Average Atomic Mass: Weighted average of isotopic masses based on natural abundance.

  • Formula:

Periodic Table

• Groups: Vertical columns; elements in the same group have similar chemical properties.

• General Classifications: Metals, nonmetals, metalloids.

• Major Groups: Alkali metals (Group 1), alkaline earth metals (Group 2), halogens (Group 17), noble gases (Group 18).

Molecules and Compounds

• Molecular Compounds: Composed of molecules formed by nonmetals sharing electrons (covalent bonds).

• Ionic Compounds: Composed of positive and negative ions (cations and anions) held together by electrostatic forces.

• Distinguishing Features:

  • Ionic compounds: Metal + nonmetal, high melting points, conduct electricity when molten or dissolved.

  • Molecular compounds: Nonmetal + nonmetal, lower melting points, do not conduct electricity.

Naming Inorganic Compounds

• Ionic Compounds: Name the cation (metal) first, then the anion (nonmetal with -ide ending).

• Transition Metals: Indicate the charge of the metal with Roman numerals (e.g., iron(III) chloride).

• Molecular Compounds: Use prefixes to indicate the number of each atom (e.g., CO2 is carbon dioxide).

• Writing Names and Formulas: Be able to write the chemical formula from the name and vice versa.

• Polyatomic Ions: Recognize common ions (e.g., sulfate SO42–, nitrate NO3–) and their charges.

Chapter 3: Stoichiometry—Chemical Formulas and Equations

Chemical Equations

Chemical equations represent the reactants and products in a chemical reaction. Balancing equations ensures the law of conservation of mass is obeyed.

• Reactants and Products: Reactants are substances consumed; products are substances formed.

• Balanced Equations: Same number of each atom on both sides of the equation.

• Types of Reactions: Synthesis, decomposition, single replacement, double replacement, combustion.

• Formula and Molecular Masses: Sum of atomic masses in a formula unit or molecule.

  • Formula mass (ionic compounds), molecular mass (molecular compounds).

• The Mole Concept: 1 mole = entities (Avogadro's number).

• Percent Composition: Percentage by mass of each element in a compound.

  • Formula:

• Empirical and Molecular Formulas:

  • Empirical formula: Simplest whole-number ratio of atoms.

  • Molecular formula: Actual number of atoms; a multiple of the empirical formula.

  • To determine molecular formula: , where

Stoichiometric Calculations

Stoichiometry involves quantitative relationships between reactants and products in a chemical reaction.

• Mole-to-Mole Problems: Use coefficients from balanced equations to relate moles of reactants and products.

• Mass-to-Mass Problems: Convert mass to moles, use mole ratio, then convert back to mass.

• Mass-to-Mole and Mole-to-Mass Problems: Use molar mass as a conversion factor.

• Example: How many grams of H2O are produced from 4.00 g of H2 and excess O2?

  • Step 1: Convert g H2 to mol H2

  • Step 2: Use mole ratio from balanced equation

  • Step 3: Convert mol H2O to g H2O

Limiting Reactant and Yield Calculations

• Limiting Reactant: The reactant that is completely consumed first, limiting the amount of product formed.

• Excess Reactant: The reactant that remains after the reaction is complete.

• Theoretical Yield: Maximum amount of product that can be formed from given reactants.

• Percent Yield:

• Example: If the theoretical yield is 10.0 g and the actual yield is 8.5 g, percent yield is .