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: Formed by combining base units (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, making it easier to express very large or small quantities.

• Kilo- (k):

• Centi- (c):

• Milli- (m):

• Micro- (\mu):

• Nano- (n):

Conversion Factors: Ratios used to convert between units (e.g., ).

Significant Figures

Significant figures reflect the precision of a measured quantity. Calculations should not imply greater precision than the measurements allow.

• 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.

• Multiplication/Division: Result has as many significant figures as the measurement with the fewest significant figures.

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

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.

• Setup: Multiply by conversion factors so units cancel appropriately.

Example: Convert 25.0 cm to meters:

Density Calculations

Density is the ratio of mass to volume and is a key property for identifying substances.

• Formula:

• Common units: g/cm3 or kg/m3

Percent Calculations

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

• Formula:

Exact vs Measured Numbers

Exact numbers have no uncertainty (e.g., counting numbers, defined conversions), while measured numbers have uncertainty and are subject to significant figure rules.

• Exact: 1 dozen = 12 items (no uncertainty)

• Measured: 2.54 cm (measured with a ruler, subject to uncertainty)

Chapter 2: Atoms, Elements, and the Periodic Table

Classification of Matter

Matter can be classified by its physical state and composition.

• Pure Substances: Elements and compounds with fixed composition.

• Mixtures: Physical combinations of two or more substances; can be homogeneous (uniform) or heterogeneous (non-uniform).

Subatomic Particles

Atoms are composed of three main subatomic particles:

• Protons: Positively charged, found in the nucleus.

• Neutrons: Neutral, found in the nucleus.

• Electrons: Negatively charged, found outside the nucleus.

Atomic Number, Mass Number, Isotopes

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

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

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

Example: Carbon-12 and Carbon-14 are isotopes of carbon.

Average Atomic Mass

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

• Formula:

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): Horizontal rows indicating energy levels.

Ions and Charges

Ions are atoms or molecules with a net electric charge due to loss or gain of electrons.

• Cations: Positively charged (loss of electrons).

• Anions: Negatively charged (gain of electrons).

Naming Compounds

Chemical nomenclature follows systematic rules for naming ionic and molecular compounds.

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

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

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: Molar mass of H2O =

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:

Percent Composition

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

• Formula:

Empirical and Molecular Formulas

The empirical formula gives the simplest whole-number ratio of elements in a compound, while the molecular formula gives the actual number of atoms of each element.

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

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

• Relationship: , where is an integer.

Example: If the empirical formula is CH2O and the molar mass is 180 g/mol, then , so the molecular formula is C6H12O6.

Additional info: These notes are based on the provided study guide and expanded with standard General Chemistry academic context to ensure completeness for exam preparation.