General Chemistry I: Fundamental Concepts and Skills

Ch.1: Introduction to Chemistry and Measurement

Key Terms and Scientific Method

This section introduces foundational chemistry vocabulary and the scientific method, which is essential for experimental design and analysis.

• Atoms, molecules, matter, state, pure substance, mixture, element, compound: Atoms are the basic units of matter. Molecules are combinations of atoms. Matter exists in different states (solid, liquid, gas). Pure substances have uniform composition (elements or compounds), while mixtures contain two or more substances physically combined. Homogeneous mixtures are uniform throughout; heterogeneous mixtures are not.

• Scientific method: Involves making observations, forming hypotheses, conducting experiments, and developing theories or laws.

• Difference between chemical and physical changes: Chemical changes result in new substances (e.g., rusting iron), while physical changes do not alter the chemical identity (e.g., melting ice).

Physical and Chemical Properties

Properties of matter are classified as physical or chemical, and can be intensive or extensive.

• Physical properties: Can be observed without changing the substance's identity (e.g., melting point, density).

• Chemical properties: Describe a substance's ability to undergo chemical changes (e.g., flammability).

• Intensive properties: Do not depend on amount (e.g., density, boiling point).

• Extensive properties: Depend on amount (e.g., mass, volume).

Measurement and Units

Understanding and converting between units is fundamental in chemistry.

• SI base units: Include meter (length), kilogram (mass), second (time), kelvin (temperature), mole (amount of substance).

• Temperature scales: Fahrenheit, Celsius, Kelvin. Conversion formulas: $T_{K} = T_{^bC} + 273.15$ $T_{^bF} = \frac{9}{5}T_{^bC} + 32$

• Metric prefixes: kilo (k), centi (c), milli (m), nano (n), etc. Example: 1 kilometer (km) = 1000 meters (m).

Accuracy, Precision, and Significant Figures

Proper measurement and reporting are crucial for scientific communication.

• Accuracy: Closeness to the true value.

• Precision: Reproducibility of measurements.

• Significant figures: Digits that carry meaning in a measurement. Rules for addition/subtraction: keep the least number of decimal places. Rules for multiplication/division: keep the least number of significant figures.

Dimensional Analysis

Dimensional analysis (factor-label method) is used to convert between units and solve problems.

• Set up conversion factors so units cancel appropriately.

• Example: Convert 25.0 cm to meters: $25.0\ \text{cm} \times \frac{1\ \text{m}}{100\ \text{cm}} = 0.250\ \text{m}$

Ch.2: Atomic Theory and the Periodic Table

Laws of Atomic Theory

Atomic theory explains the nature of matter and the relationships between elements and compounds.

• Law of conservation of mass: Mass is neither created nor destroyed in a chemical reaction.

• Law of definite proportions: A compound always contains the same elements in the same proportion by mass.

• Law of multiple proportions: When two elements form more than one compound, the ratios of the masses of the second element that combine with a fixed mass of the first element are ratios of small whole numbers.

Structure of the Atom

Atoms consist of 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.

• Atomic number (Z): Number of protons.

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

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

• Average atomic mass: Weighted average of all isotopes.

The Periodic Table

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

• Groups (columns): Elements with similar properties (e.g., alkali metals, halogens).

• Periods (rows): Horizontal rows showing trends in properties.

• Metals, nonmetals, metalloids: Classified by physical and chemical properties.

• Transition metals: Elements in the d-block, often form colored compounds and have variable oxidation states.

Chemical Bonding and Compounds

Ionic and Covalent Bonding

Chemical bonds form between atoms to achieve stability. The two main types are ionic and covalent bonds.

• Ionic bonding: Transfer of electrons from a metal to a nonmetal, forming ions.

• Covalent bonding: Sharing of electrons between nonmetals.

• Identifying bond type: Use the formula and periodic table to determine if a compound is ionic or covalent.

Naming Compounds and Polyatomic Ions

Systematic naming allows chemists to communicate compound identities unambiguously.

• Ionic compounds: Name the cation first, then the anion. Include Roman numerals for transition metals with variable charge.

• Polyatomic ions: Common examples include phosphate (PO43-), sulfate (SO42-), nitrate (NO3-), carbonate (CO32-), ammonium (NH4+), perchlorate (ClO4-), cyanide (CN-), nitrite (NO2-), hydroxide (OH-).

• Covalent compounds: Use prefixes (mono-, di-, tri-, etc.) to indicate the number of each atom.

Formulas and Representations

Compounds can be represented in several ways to convey structure and composition.

• Empirical formula: Simplest whole-number ratio of atoms.

• Molecular formula: Actual number of atoms of each element.

• Structural formula: Shows how atoms are connected.

• Models: Ball-and-stick, space-filling, and Lewis structures.

Stoichiometry and Chemical Calculations

Mole Concept and Molar Mass

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

• Mole (mol): $6.022 \times 10^{23}$ particles (Avogadro's number).

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

• Converting mass to moles: $\text{moles} = \frac{\text{mass (g)}}{\text{molar mass (g/mol)}}$

• Converting moles to number of particles: $\text{particles} = \text{moles} \times 6.022 \times 10^{23}$

Percent Composition and Empirical/Molecular Formulas

Percent composition describes the mass percentage of each element in a compound. Empirical and molecular formulas are determined from composition data.

• Percent composition: $\%\ \text{element} = \frac{\text{mass of element in 1 mol compound}}{\text{molar mass of compound}} \times 100\%$

• Empirical formula: Simplest ratio of elements, found by dividing moles of each element by the smallest number of moles.

• Molecular formula: Multiple of the empirical formula, determined by dividing the compound's molar mass by the empirical formula mass.

Sample Table: Common Polyatomic Ions

This table summarizes the names, formulas, and charges of important polyatomic ions.

Name	Formula	Charge
Phosphate	PO4	3-
Sulfate	SO4	2-
Nitrate	NO3	1-
Carbonate	CO3	2-
Ammonium	NH4	1+
Perchlorate	ClO4	1-
Cyanide	CN	1-
Nitrite	NO2	1-
Hydroxide	OH	1-

Using Formulas in Calculations

• Use molar mass as a conversion factor between grams and moles.

• Use percent composition to determine empirical and molecular formulas.

• Apply dimensional analysis to solve multi-step problems involving moles, mass, and number of particles.

Example: Calculating Empirical Formula

If a compound contains 40.0% carbon, 6.7% hydrogen, and 53.3% oxygen by mass:

• Assume 100 g sample: 40.0 g C, 6.7 g H, 53.3 g O

• Convert to moles: $\text{mol C} = \frac{40.0}{12.01} = 3.33$ $\text{mol H} = \frac{6.7}{1.008} = 6.65$ $\text{mol O} = \frac{53.3}{16.00} = 3.33$

• Divide by smallest: C: 1, H: 2, O: 1 → Empirical formula: CH2O