Chapter 1: The Chemical World

Scientific Method

The scientific method is a systematic approach used by scientists to explore observations, answer questions, and solve problems. It ensures that scientific inquiry is logical, repeatable, and based on evidence.

• Observation: Gathering information through the senses or instruments.

• Hypothesis: A testable explanation for an observation.

• Experiment: A controlled procedure to test the hypothesis.

• Analysis: Interpreting data to determine if it supports the hypothesis.

• Conclusion: A summary of findings; may lead to further hypotheses or theories.

Example: Observing that ice melts at room temperature, hypothesizing that heat causes melting, and designing an experiment to test this idea.

Chapter 2: Measurement and Problem Solving

Significant Figures in Measurements and Calculations

Significant figures reflect the precision of a measured or calculated quantity.

• 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.

Example: 0.00450 has three significant figures.

Exponential Numbers: Writing and Use in Calculations

Exponential notation (scientific notation) expresses numbers as a product of a coefficient and a power of ten.

• General form:

• Example: 3,200 =

Unit Analysis Problem Solving

Unit analysis (dimensional analysis) uses conversion factors to solve problems involving measurements.

• Set up conversion factors so units cancel appropriately.

• Example: To convert 5.0 cm to meters:

Metric Conversions

Common metric prefixes to memorize:

• Kilo- (k):

• Centi- (c):

• Milli- (m):

Volume Conversions

When converting units of volume, remember to cube both the number and the unit.

• Example:

Density in Calculations

Density is the mass per unit volume of a substance.

• Formula:

• Example: If a block has a mass of 10 g and a volume of 2 cm3, its density is .

Chapter 3: Matter and Energy

Phases of Matter and Phase Changes

Matter exists in three main phases: solid, liquid, and gas. Phase changes include:

• Melting: Solid to liquid

• Freezing: Liquid to solid

• Vaporization: Liquid to gas

• Condensation: Gas to liquid

• Sublimation: Solid to gas

Classification of Matter

• Element: Pure substance made of one type of atom (e.g., O2).

• Compound: Substance made of two or more elements chemically combined (e.g., H2O).

• Homogeneous mixture: Uniform composition (e.g., saltwater).

• Heterogeneous mixture: Non-uniform composition (e.g., salad).

Physical and Chemical Properties and Changes

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

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

• Physical change: Does not alter composition (e.g., melting ice).

• Chemical change: Alters composition (e.g., rusting iron).

Conservation of Mass

Mass is neither created nor destroyed in a chemical reaction.

• Law of Conservation of Mass: Total mass of reactants equals total mass of products.

Energy, Exothermic, Endothermic, Heat Capacity

• Energy: The capacity to do work or produce heat.

• Exothermic process: Releases energy (e.g., combustion).

• Endothermic process: Absorbs energy (e.g., melting ice).

• Heat capacity (C): Amount of heat needed to raise temperature by 1°C.

• Formula:

Temperature

Temperature measures the average kinetic energy of particles. Common scales: Celsius (°C), Kelvin (K), Fahrenheit (°F).

Additional info: Temperature conversions are not required to be memorized for this course.

Chapter 4: Atoms and Elements

Dalton Model of the Atom

John Dalton proposed that matter is made of indivisible atoms, each element consisting of identical atoms.

Subatomic Particles

• Proton: Positive charge, located in nucleus, mass ≈ 1 amu.

• Neutron: No charge, located in nucleus, mass ≈ 1 amu.

• Electron: Negative charge, orbits nucleus, mass ≈ 1/1836 amu.

Atomic Notation

Represents the number of protons, neutrons, and electrons in an atom.

• General form: , where X = element symbol, A = mass number, Z = atomic number.

Isotopes

Atoms of the same element with different numbers of neutrons.

• Example: and are isotopes of carbon.

Mass Number and Atomic Mass

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

• Atomic mass: Weighted average mass of all isotopes (in amu).

Periodic Law and Classification of Elements

The periodic law states that properties of elements repeat periodically when arranged by atomic number.

• Metals: Shiny, conductive, malleable (e.g., Fe, Cu).

• Nonmetals: Dull, poor conductors (e.g., O, N).

• Metalloids: Properties intermediate between metals and nonmetals (e.g., Si).

• Main group elements: Groups 1, 2, 13-18.

• Transition metals: Groups 3-12.

• Rare earth elements: Lanthanides and actinides.

• Alkali metals: Group 1 (e.g., Na, K).

• Alkaline earth metals: Group 2 (e.g., Mg, Ca).

• Halogens: Group 17 (e.g., F, Cl).

• Noble gases: Group 18 (e.g., He, Ne).

Chapter 5: Molecules and Compounds

Compounds and Chemical Formulas

A compound is a substance composed of two or more elements in fixed ratios, represented by a chemical formula (e.g., H2O).

Naming Ionic, Molecular, and Acid Compounds

• Ionic compounds: Metal + nonmetal; name cation first, then anion (e.g., NaCl: sodium chloride).

• Molecular compounds: Nonmetal + nonmetal; use prefixes (e.g., CO2: carbon dioxide).

• Acids: Compounds that release H+ in water; naming depends on anion (e.g., HCl: hydrochloric acid).

Chapter 6: Chemical Composition

The Mole

The mole is the SI unit for amount of substance. One mole contains Avogadro's number () of particles.

Molar Mass Calculations

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

• Calculate by summing atomic masses from the periodic table.

• Example: Molar mass of H2O = g/mol.

Conversions: Grams, Moles, and Number of Particles

• Use the following relationships:

Percent Composition

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

• Formula:

Empirical and Molecular Formulas

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

• Molecular formula: Actual number of atoms in a molecule.

• To find empirical formula: Convert masses or percentages to moles, divide by smallest, write ratio.

• To find molecular formula: , where

Chapter 9: Electrons in Atoms and the Periodic Table

Energy Levels and Sublevels

Electrons occupy energy levels (shells) and sublevels (s, p, d, f) around the nucleus.

• Principal energy levels: n = 1, 2, 3, ...

• Sublevels: s (2 electrons), p (6), d (10), f (14)

Electron Configurations

Electron configuration shows the arrangement of electrons in an atom or ion.

• Follow the Aufbau principle, Pauli exclusion principle, and Hund's rule.

• Example: Oxygen (O): 1s2 2s2 2p4

Octet Rule and Ionic Compound Formation

Atoms tend to gain, lose, or share electrons to achieve eight valence electrons (octet rule).

• Metals lose electrons to form cations; nonmetals gain electrons to form anions.

• Ionic compounds form from electrostatic attraction between cations and anions.

Periodic Trends

• Atomic radius: Increases down a group, decreases across a period.

• Ionization energy: Energy to remove an electron; decreases down a group, increases across a period.

• Metallic character: Increases down a group, decreases across a period.

Elements to Memorize

Students should memorize the names and symbols of the following elements:

• H, He, Li, Be, B, C, N, O, F, Ne, Na, Mg, Al, Si, P, S, Cl, Ar, K, Ca, Cr, Mn, Fe, Co, Ni, Cu, Zn, Se, Br, Kr, Ag, Sn, I, Xe, Ba, Pt, Au, Hg, Pb, U

Practice Midterm

Completing the practice midterm is required for exam preparation.