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, joule (J) for energy ().

• Example: The unit for density is kilograms per cubic meter ().

Metric Prefixes and Conversion Factors

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

• Common Prefixes: kilo- (k, ), centi- (c, ), milli- (m, ), micro- (, ), nano- (n, ).

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

• Example: To convert 2500 mg to grams: .

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 5.0 miles to kilometers: .

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 10.0 g and a volume of 2.0 cm3, its density is .

Percent Calculations

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

• Formula:

• Example: If 25 g of NaCl is dissolved in 100 g of water, the percent by mass of NaCl is .

Exact vs Measured Numbers

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

• Exact Numbers: Defined values or counted quantities (e.g., 1 dozen = 12, 1000 m = 1 km) with infinite significant figures.

• Measured Numbers: Obtained by measurement, subject to uncertainty, and limited by the precision of the instrument.

• Example: The number of students in a classroom (exact); the mass of a sample measured on a balance (measured).

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 fixed composition.

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

• Example: Air (homogeneous mixture), sand and iron filings (heterogeneous mixture).

Subatomic Particles

Atoms are composed of three primary subatomic particles.

• Protons: Positively charged, located in the nucleus, mass ≈ 1 amu.

• Neutrons: Neutral, located in the nucleus, mass ≈ 1 amu.

• Electrons: Negatively charged, located outside the nucleus, mass ≈ 0.0005 amu.

• 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 amu) and (24.2%, 36.97 amu): amu.

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 loses one electron to form ; Cl gains one electron to form .

Naming Compounds

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

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

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

• Example: is calcium chloride; is dinitrogen pentoxide.

Atomic Theory Contributors

Key scientists contributed to the development of atomic theory.

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

• J.J. Thomson: Discovered the electron (cathode ray experiment).

• Ernest Rutherford: Discovered the nucleus (gold foil experiment).

• James Chadwick: Discovered the neutron.

• Example: Rutherford's experiment led to the nuclear model of the atom.

Chapter 3: Molecules, Compounds, and the Mole

Molar Mass

The 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: : 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 to moles:

  • Moles to particles:

• Example: 36.0 g of water: mol.

Percent Composition

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

• Formula:

• Example: In , percent H: .

Empirical and Molecular Formulas

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

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

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

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

• Example: If empirical formula is CH2O (30 g/mol) and molar mass is 180 g/mol, molecular formula is C6H12O6.

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.