Chapter 1: Matter, Measurement, and Problem Solving

SI Base Units and Derived Units

The International System of Units (SI) provides standard units for scientific measurements. Understanding these units is essential for accurate communication and calculation in chemistry.

• SI Base Units: Fundamental units for physical quantities (e.g., meter for length, kilogram for mass, second for time).

• Derived Units: Combinations of base units to express other quantities (e.g., meters per second for speed, cubic meters for volume).

• Example: The SI unit for density is kilograms per cubic meter (kg/m3).

Metric Prefixes and Conversion Factors

Metric prefixes indicate multiples or fractions of base units, facilitating conversions between different scales.

• Common Prefixes: kilo- (103), centi- (10-2), milli- (10-3), micro- (10-6), nano- (10-9).

• Conversion Factors: Ratios used to convert between units (e.g., 1 km = 1000 m).

• Example: To convert 5.0 cm to meters:

Significant Figures

Significant figures reflect the precision of a measured quantity. Proper use is crucial for reporting and calculating scientific data.

• Rules: All nonzero digits are significant; zeros between nonzero digits are significant; leading zeros are not significant; trailing zeros are significant if after a decimal point.

• Calculations: For multiplication/division, the result has as many significant figures as the measurement with the fewest. For addition/subtraction, the result has as many decimal places as the measurement with the fewest decimal places.

• Example: (rounded to two significant figures)

Dimensional Analysis

Dimensional analysis (factor-label method) is a systematic approach to problem solving that uses conversion factors to move from one unit to another.

• Process: Multiply by conversion factors so that units cancel appropriately.

• Example: Convert 25.0 inches to centimeters:

Density Calculations

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

• Formula:

• Units: Commonly g/cm3 or kg/m3.

• Example: If a block has a mass of 50.0 g and a volume of 20.0 cm3, its density is

Percent Calculations

Percentages are used to express proportions, such as percent composition in compounds.

• Formula:

• Example: If 5 g of salt is dissolved in 100 g of water, the percent by mass is

Exact vs Measured Numbers

Understanding the difference between exact and measured numbers is important for determining significant figures.

• Exact Numbers: Defined values or counts (e.g., 1 dozen = 12, 100 cm = 1 m) with infinite significant figures.

• Measured Numbers: Obtained by measurement, limited by instrument precision, and have a specific number of significant figures.

Chapter 2: Atoms, Elements, and the Periodic Table

Classification of Matter

Matter can be classified based on its physical and chemical properties.

• Pure Substances: Elements and compounds with uniform composition.

• Mixtures: Physical combinations of substances; can be homogeneous (solutions) or heterogeneous.

• Example: Air is a homogeneous mixture; sand and iron filings form a heterogeneous mixture.

Subatomic Particles

Atoms are composed of three primary subatomic particles.

• Protons: Positively charged, located in the nucleus.

• Neutrons: Neutral, located in the nucleus.

• Electrons: Negatively charged, found in orbitals around the nucleus.

• Example: A carbon atom has 6 protons, 6 neutrons, and 6 electrons.

Atomic Number, Mass Number, Isotopes

Atoms of the same element can have different numbers of neutrons, resulting in isotopes.

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

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

• Isotopes: Atoms with the same atomic number but different mass numbers.

• Example: and are isotopes of carbon.

Average Atomic Mass

The average atomic mass of an element reflects the weighted average of all naturally occurring isotopes.

• Formula:

• Example: If (75.8%, 34.97 u) and (24.2%, 36.97 u): u

Periodic Table Groups

The periodic table organizes elements by increasing atomic number and similar chemical properties.

• Groups (Columns): Elements with similar properties (e.g., Group 1: Alkali metals, Group 17: Halogens).

• Periods (Rows): Indicate energy levels of electrons.

Ions and Charges

Atoms can gain or lose electrons to form ions.

• Cations: Positively charged ions (loss of electrons).

• Anions: Negatively charged ions (gain of electrons).

• Example: Na → Na+ + e-; Cl + e- → Cl-

Naming Compounds

Chemical nomenclature follows systematic rules for naming compounds.

• Ionic Compounds: Name cation first, then anion (e.g., NaCl: sodium chloride).

• Covalent Compounds: Use prefixes to indicate number of atoms (e.g., CO2: carbon dioxide).

• Example: FeCl3 is iron(III) chloride.

Atomic Theory Contributors

Key scientists contributed to the development of atomic theory.

• John Dalton: Proposed the atomic theory (elements are made of atoms).

• J.J. Thomson: Discovered the electron.

• Ernest Rutherford: Discovered the nucleus.

• James Chadwick: Discovered the neutron.

Chapter 3: Molecules, Compounds, and the Mole

Molar Mass

Molar mass is the mass of one mole of a substance, expressed in grams per mole (g/mol).

• Calculation: Sum the atomic masses of all atoms in a formula.

• Example: H2O: (2 × 1.01) + 16.00 = 18.02 g/mol

Mole Conversions

The mole is a counting unit in chemistry, relating mass, number of particles, and volume (for gases).

• Avogadro's Number: particles per mole.

• Conversions:

  • Mass ↔ Moles:

  • Moles ↔ Particles:

• Example: 36.0 g H2O × = 2.0 mol H2O

Percent Composition

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

• Formula:

• Example: In H2O, percent H =

Empirical and Molecular Formulas

The empirical formula gives the simplest whole-number ratio of atoms, while the molecular formula gives the actual number of atoms in a molecule.

• Empirical Formula: Simplest ratio (e.g., CH2O for glucose).

• Molecular Formula: Actual number (e.g., C6H12O6 for glucose).

• Calculation: Divide the molecular mass by the empirical formula mass to find the multiple.

• Example: If empirical formula mass = 30 g/mol, molecular mass = 60 g/mol, then molecular formula is twice the empirical formula.

Additional info: This guide covers foundational concepts for General Chemistry, including measurement, atomic structure, and basic chemical calculations. Students should also review assigned readings and practice problems for comprehensive preparation.