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, such as meter (m) for length, kilogram (kg) for mass, second (s) for time, mole (mol) for amount of substance, kelvin (K) for temperature, ampere (A) for electric current, and candela (cd) for luminous intensity.

• Derived Units: Formed by combining base units, e.g., meters per second (m/s) for speed, cubic meters (m3) for volume.

• Example: The 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- (k, 103), centi- (c, 10-2), milli- (m, 10-3), micro- (μ, 10-6), nano- (n, 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 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 units cancel appropriately.

• Example: Convert 25.0 inches to centimeters:

Density Calculations

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

• 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,

Percent Calculations

Percent calculations are used to express the ratio of a part to a whole as a percentage.

• Formula:

• Example: If 5 g of salt is dissolved in 100 g of water,

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 two or more 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 main subatomic particles: protons, neutrons, and electrons.

• Protons: Positively charged, found in the nucleus.

• Neutrons: Neutral, found 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

Each element is defined by its atomic number, while isotopes differ in mass number.

• 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 with different numbers of neutrons.

• 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 chlorine has two isotopes, (75.8%, 34.97 u) and (24.2%, 36.97 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.

• Example: Sodium (Na) is in Group 1, Period 3.

Ions and Charges

Atoms can gain or lose electrons to form ions, which carry a net charge.

• 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/mol.

• Conversions:

  • Mass ↔ Moles:

  • Moles ↔ Particles:

• Example: 36.0 g H2O:

Percent Composition

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

• Formula:

• Example: In H2O, %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:

  • Find moles of each element.

  • Divide by smallest number of moles.

  • Multiply to get whole numbers if necessary.

• Example: A compound with 40.0% C, 6.7% H, 53.3% O has empirical formula CH2O.

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.