General Chemistry Exam I Study Guide

Significant Figures and Measurement

Understanding significant figures is essential for reporting scientific measurements accurately. Significant figures reflect the precision of a measured or calculated value.

• Significant Figures: The digits in a measurement that are known with certainty plus one estimated digit.

• Rules:

  • All nonzero digits are significant.

  • Zeros between nonzero digits are significant.

  • Leading zeros are not significant.

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

• Reporting Calculated Values: Round the final answer to the correct number of significant figures based on the operation performed (addition/subtraction: least decimal places; multiplication/division: least significant figures).

• Example: (rounded to 2 significant figures)

Density and Unit Conversions

Density is a fundamental property that relates the mass and volume of a substance.

• Definition: Density () is mass per unit volume.

• Formula:

• Units: Commonly expressed in or for solids and liquids, for gases.

• Unit Conversions: Use dimensional analysis to convert between units (e.g., ).

• Example: If a metal has a mass of and displaces of water,

Matter: Classification and Properties

Matter can be classified based on its composition and uniformity.

• Pure Substances: Have a fixed composition; can be elements or compounds.

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

• Physical vs. Chemical Changes:

  • Physical Change: Does not alter the chemical composition (e.g., melting, boiling).

  • Chemical Change: Produces new substances (e.g., combustion, rusting).

• Example: Salt water is a homogeneous mixture; sand and iron filings is a heterogeneous mixture.

Atoms, Isotopes, and Ions

Atoms are the basic units of matter, composed of protons, neutrons, and electrons.

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

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

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

• Ions: Atoms or molecules with a net charge due to loss or gain of electrons.

• Determining Subatomic Particles:

  • Protons = Atomic number ()

  • Neutrons = Mass number () - Atomic number ()

  • Electrons = Protons (for neutral atoms); adjust for charge in ions

• Example: has 11 protons, 12 neutrons, and 10 electrons.

Naming Compounds and Chemical Formulas

Chemical nomenclature allows chemists to communicate the composition of compounds unambiguously.

• Ionic Compounds: Composed of metals and nonmetals; named using cation and anion names.

• Type I: Metal forms only one type of cation (e.g., NaCl: sodium chloride).

• Type II: Metal forms more than one cation; use Roman numerals (e.g., FeCl2: iron(II) chloride).

• Covalent Compounds: Composed of nonmetals; use prefixes (e.g., CO2: carbon dioxide).

• Hydrocarbons: Compounds of carbon and hydrogen; types of bonds (single, double, triple) determine the class (alkanes, alkenes, alkynes).

• Example: is ethane (an alkane with single bonds).

Empirical and Molar 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.

• Empirical Formula: Simplest ratio of elements.

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

• Calculating Empirical Formula:

  1. Convert mass percentages to grams (assume 100 g sample).

  2. Convert grams to moles for each element.

  3. Divide by the smallest number of moles to get the ratio.

• Example: A compound with 40% C, 6.7% H, and 53.3% O has empirical formula CH2O.

Physical and Chemical Properties

Properties of matter can be classified as physical or chemical.

• Physical Properties: Observed without changing composition (e.g., melting point, density).

• Chemical Properties: Observed during a chemical change (e.g., flammability, reactivity).

• Example: Boiling water is a physical change; burning wood is a chemical change.

Balancing Chemical Equations

Balanced chemical equations reflect the conservation of mass and atoms.

• Steps:

  1. Write the unbalanced equation.

  2. Count atoms of each element on both sides.

  3. Add coefficients to balance atoms.

• Example:

Temperature Conversions

Temperature can be measured in Celsius, Kelvin, or Fahrenheit. Chemistry commonly uses Celsius and Kelvin.

• Conversion Formulas:

• Example:

Periodic Table and Element Classification

The periodic table organizes elements by atomic number and properties.

• Groups: Vertical columns; elements in a group have similar properties.

• Periods: Horizontal rows.

• Main Group Elements: Groups 1, 2, and 13-18.

• Transition Metals: Groups 3-12.

• Special Groups:

  • Alkali Metals: Group 1

  • Alkaline Earth Metals: Group 2

  • Halogens: Group 17

  • Noble Gases: Group 18

• Example: Sodium (Na) is an alkali metal; chlorine (Cl) is a halogen.

Dalton’s Atomic Theory

Dalton’s atomic theory laid the foundation for modern chemistry.

• All matter is composed of indivisible atoms.

• Atoms of the same element are identical; atoms of different elements are different.

• Atoms combine in simple whole-number ratios to form compounds.

• Atoms are rearranged in chemical reactions, not created or destroyed.

• Note: Some parts of Dalton’s theory have been revised (e.g., atoms are divisible into subatomic particles).

Rutherford’s Nuclear Model

Rutherford’s gold foil experiment led to the nuclear model of the atom.

• Most of the atom’s mass and all positive charge are concentrated in a small nucleus.

• Electrons occupy the space around the nucleus.

• The atom is mostly empty space.

Bonding: Ionic, Covalent, and Polar Covalent

Chemical bonds form when atoms share or transfer electrons.

• Ionic Bonds: Transfer of electrons from metal to nonmetal; forms ions (e.g., NaCl).

• Covalent Bonds: Sharing of electrons between nonmetals (e.g., H2O).

• Polar Covalent Bonds: Unequal sharing of electrons between different nonmetals (e.g., HCl).

• Example: In NaCl, Na donates an electron to Cl, forming Na+ and Cl-.

Calculations Involving Moles, Mass, and Atoms

The mole is a counting unit in chemistry, relating mass to number of particles.

• Avogadro’s Number: particles/mol.

• Molar Mass: Mass of one mole of a substance (g/mol).

• Conversions:

  • Mass to moles:

  • Moles to atoms:

• Example: of C () contains mole or atoms.

Properties of Metals and Nonmetals

Elements are classified as metals, nonmetals, or metalloids based on their properties.

• Metals: Shiny, good conductors, malleable, ductile.

• Nonmetals: Dull, poor conductors, brittle.

• Metalloids: Properties intermediate between metals and nonmetals.

• Example: Iron is a metal; sulfur is a nonmetal; silicon is a metalloid.

Average Atomic Mass

The average atomic mass of an element reflects the weighted average of its isotopes.

• Formula:

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

Volume and Unit Comparisons

Comparing quantities with different units requires conversion to a common unit.

• Common Volume Units:

• Example: is larger than .

Sample Table: Classification of Matter

The following table summarizes the classification of matter:

Additional info:

• Some content (e.g., the 'mystery extra question') is unspecified and not included.

• For exam preparation, practice problems involving calculations, naming, and classification are recommended.