Chapter 1: Chemistry in Our Lives

What is Chemistry?

Chemistry is the study of matter, its properties, the changes it undergoes, and the energy associated with those changes. A chemical is any substance that has a definite composition.

• Matter: Anything that has mass and occupies space.

• Chemical: A substance with a specific composition (e.g., water, carbon dioxide).

The Scientific Method

The scientific method is a systematic approach to understanding the natural world through observation and experimentation.

• Hypothesis: A tentative explanation for an observation.

• Theory: A well-tested explanation for a broad set of observations.

• Law: A statement that summarizes the results of many observations and experiments.

• Steps: Observation → Hypothesis → Experiment → Analysis → Conclusion.

States of Matter

• Solid: Definite shape and volume.

• Liquid: Definite volume, takes shape of container.

• Gas: No definite shape or volume.

Basic Calculations and Data Interpretation

• Percentage Calculation:

• Solving for Variables: Rearranging equations to isolate the desired variable.

• Graph Interpretation: Understanding trends, axes, and data points.

Chapter 2: Chemistry and Measurements

SI Units and Measurement

The International System of Units (SI) is used for scientific measurements.

• Length: meter (m)

• Mass: kilogram (kg)

• Time: second (s)

• Temperature: kelvin (K)

• Amount of substance: mole (mol)

Rounding and Uncertainty

• Round up if the next digit is 5 or more; round down if less than 5.

• Measurements include all certain digits plus one uncertain digit.

Significant Figures (Sig Figs)

• Rules for zeros: Leading zeros are not significant; captive zeros are significant; trailing zeros are significant if there is a decimal point.

• Multiplication/Division: Result has as many sig figs as the measurement with the fewest sig figs.

• Addition/Subtraction: Result has as many decimal places as the measurement with the fewest decimal places.

• Exact Numbers: Numbers from counting or defined values (e.g., 1 in = 2.54 cm) do not limit sig figs.

Scientific Notation

• Expresses numbers as a product of a coefficient and a power of 10.

• Example:

Metric Prefixes and Equalities

• Common prefixes: kilo- (k, ), centi- (c, ), milli- (m, ), micro- (, )

• Equalities: 1 kg = 2.205 lbs; 1 in = 2.54 cm (exact)

Conversions and Dimensional Analysis

• Use conversion factors to change units.

• Set up problems so units cancel appropriately.

• Example:

Density and Specific Gravity

• Density: (units: g/mL or g/cm3)

• Specific Gravity: (unitless)

• Typical value for urine: 1.003–1.030

Chapter 3: Matter and Energy

Classification of Matter

• Element: Pure substance, cannot be broken down (e.g., O2).

• Compound: Pure substance, can be broken down into elements (e.g., H2O).

• Mixture: Physical blend of two or more substances.

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

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

Separation of Mixtures

• Physical methods: Filtration, distillation, chromatography.

Physical vs. Chemical Properties and Changes

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

• Chemical property: Describes ability to undergo change (e.g., flammability).

• Physical change: Change in state or appearance, not composition (e.g., melting ice).

• Chemical change: Produces new substances (e.g., rusting iron).

Temperature and Energy

• Temperature conversions (formulas provided on exam):

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

• Heat: Energy transferred due to temperature difference.

• Work: Energy used to move an object against a force.

• Kinetic energy: Energy of motion.

• Potential energy: Stored energy.

Units of Energy

• calorie (cal): Energy to raise 1 g of water by 1°C.

• kilocalorie (kcal): 1,000 cal.

• joule (J): SI unit of energy.

• kilojoule (kJ): 1,000 J.

• 1 cal = 4.184 J

Energy Values of Foods

• Carbohydrates: 4 Cal/g

• Proteins: 4 Cal/g

• Fats: 9 Cal/g

• Total Calories: Multiply grams by Cal/g for each nutrient.

Specific Heat

• Formula:

• q = heat (J), m = mass (g), c = specific heat (J/g·°C), = temperature change (°C)

Changes of State

• Melting (solid to liquid), Freezing (liquid to solid)

• Vaporization (liquid to gas), Condensation (gas to liquid)

• Sublimation (solid to gas), Deposition (gas to solid)

• Heat of fusion: Energy to melt 1 g of substance.

• Heat of vaporization: Energy to vaporize 1 g of substance.

• Heating curve: Graph showing temperature change as heat is added.

Chapter 4: Atoms and Elements

Elements and the Periodic Table

• Know names and symbols for elements 1–20.

• Table 4.2: (Additional info: Table likely lists element names, symbols, and atomic numbers for 1–20.)

• Groups: Vertical columns; Periods: Horizontal rows.

• Representative elements: Groups 1A–8A.

• Transition elements: Groups 3–12.

• Metals: Left and center; Nonmetals: Right; Metalloids: Border between metals and nonmetals.

• Group names:

  • 1A: Alkali metals

  • 2A: Alkaline earth metals

  • 7A: Halogens

  • 8A: Noble gases

Atoms and Atomic Theory

• Atom: Smallest unit of an element.

• Molecule: Two or more atoms bonded together.

• Dalton’s Atomic Theory (4 postulates):

  1. All matter is made of atoms.

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

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

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

• Rutherford’s Gold Foil Experiment: Discovered nucleus; atoms are mostly empty space; nucleus is dense and positively charged.

Structure of the Atom

• Atomic number (Z): Number of protons.

• Mass number (A): Protons + neutrons.

• Proton: +1 charge, mass ≈ 1 amu.

• Neutron: 0 charge, mass ≈ 1 amu.

• Electron: –1 charge, mass ≈ 0.0005 amu.

• Atomic mass unit (amu): 1/12 the mass of a carbon-12 atom.

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

• Calculating neutrons:

• Average atomic mass: Weighted average of isotopes’ masses.

• Atomic (nuclear) notation: (X = element symbol)

Electron Arrangement

• Electrons occupy energy levels (shells).

• Maximum electrons per level: 2, 8, 18, 32…

• Core electrons: Inner shells; Valence electrons: Outermost shell.

• Valence shell: Highest energy level with electrons.

• Lewis symbols: Dots around element symbol representing valence electrons.

Periodic Trends

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

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

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

Table: Properties of Subatomic Particles

Additional info: Table 4.2 likely lists the first 20 elements with their names, symbols, and atomic numbers. Students should memorize these for the exam.