GOB Chemistry Study Guide: Chapters 1–5 (Chemistry in Our Lives, Measurements, Matter & Energy, Atoms & Elements, Nuclear Chemistry)

Ch 1: Chemistry in Our Lives

1.1 Chemistry and Chemicals

Chemistry is the science that studies the composition, structure, properties, and reactions of matter. A chemical is a substance with a definite composition and set of properties, regardless of its source.

• Chemistry: The study of matter and its changes.

• Chemical: Always has the same composition and properties wherever found.

• Example: Water (H2O) is a chemical whether found in a river or a laboratory.

1.2 The Scientific Method

The scientific method is a systematic approach to understanding natural phenomena.

• Begins with observations.

• Form a hypothesis (a tentative explanation).

• Perform experiments to test the hypothesis.

• If experiments support the hypothesis, a conclusion is drawn.

1.3 Strategies for Learning Chemistry

Effective learning in chemistry involves active engagement with the material and structured study plans.

• Use textbook features: Learning Goals, Engage questions, Sample Problems, Study Checks, Practice Problems, Chapter Review, Key Terms, and Answers.

• Reflect on your understanding with features like "Does Your Answer Make Sense?" and "Thinking About Your Thinking."

1.4 Math Concepts in Chemistry

Basic math skills are essential for solving chemistry problems.

• Understanding place values, positive and negative numbers, percentages, solving equations, and interpreting graphs.

1.5 Scientific Notation

Scientific notation expresses numbers as a product of a coefficient and a power of ten.

• General form: $a \times 10^n$

• If the number is greater than 1, $n$ is positive; if less than 1, $n$ is negative.

• Example: 4,500 = $4.5 \times 10^3$; 0.0062 = $6.2 \times 10^{-3}$

Ch 2: Chemistry & Measurements

2.1 Units of Measurement

Measurements in science use the metric system or SI (International System of Units).

• Volume: liter (L)

• Length: meter (m)

• Mass: gram (g), kilogram (kg)

• Temperature: degree Celsius ($^\circ$C), kelvin (K)

• Time: second (s)

2.2 Measured vs. Exact Numbers & Significant Figures

• Measured numbers: Obtained using measuring devices; have uncertainty.

• Exact numbers: Obtained by counting or definition; no uncertainty.

• Significant figures: All certain digits plus one estimated digit in a measurement.

• Leading zeros (before decimal) and trailing zeros (after non-decimal) are not significant.

2.3 Significant Figures in Calculations

• Multiplication/Division: Result has same number of significant figures as the measurement with the fewest significant figures.

• Addition/Subtraction: Result has same number of decimal places as the measurement with the fewest decimal places.

2.4 Metric Prefixes and Equalities

Metric prefixes change the size of a unit by powers of ten.

• Smaller units: centi (c, $10^{-2}$), milli (m, $10^{-3}$), micro ($\mu$, $10^{-6}$)

• Larger units: kilo (k, $10^3$), mega (M, $10^6$), tera (T, $10^{12}$)

• Equalities: $1\ \mathrm{L} = 1000\ \mathrm{mL}$, $1\ \mathrm{m} = 100\ \mathrm{cm}$, $1\ \mathrm{kg} = 1000\ \mathrm{g}$, $1\ \mathrm{min} = 60\ \mathrm{s}$

2.5 Conversion Factors

• Express relationships as fractions to convert between units.

• Each equality gives two conversion factors (e.g., $\frac{1\ \mathrm{L}}{1000\ \mathrm{mL}}$ and $\frac{1000\ \mathrm{mL}}{1\ \mathrm{L}}$).

• Percentages can be used as conversion factors (e.g., 5% = $\frac{5}{100}$).

2.6 Unit Conversions

• Multiply the given value by conversion factors so that units cancel, leaving the desired unit.

• Example: Convert 250 mL to L: $250\ \mathrm{mL} \times \frac{1\ \mathrm{L}}{1000\ \mathrm{mL}} = 0.25\ \mathrm{L}$

2.7 Density and Specific Gravity

• Density: Ratio of mass to volume. $\text{Density} = \frac{\text{mass}}{\text{volume}}$

• Common units: g/mL or $\mathrm{g/cm^3}$

• Specific gravity: Ratio of the density of a substance to the density of water (1.000 g/mL).

• Density can be used as a conversion factor to find mass or volume.

Ch 3: Matter & Energy

3.1 Classification of Matter

• Matter: Anything with mass and volume.

• Pure substances: Elements or compounds with fixed composition.

• Mixtures: Variable composition; can be separated by physical methods.

• Example: Salt water is a mixture; sodium chloride is a pure substance.

3.2 States and Properties of Matter

• States: Solid, liquid, gas.

• Physical property: Observed/measured without changing identity (e.g., color, melting point).

• Chemical property: Ability to change into another substance (e.g., flammability).

• Physical change: No new substance formed (e.g., melting ice).

• Chemical change: New substance formed (e.g., rusting iron).

3.3 Temperature Scales and Conversions

• Celsius: Freezing point 0°C, boiling point 100°C.

• Fahrenheit: Freezing point 32°F, boiling point 212°F.

• Kelvin: SI unit; $T_K = T_C + 273$

• Conversion: $T_F = 1.8T_C + 32$

3.4 Energy: Potential and Kinetic

• Energy: Ability to do work.

• Potential energy: Stored energy.

• Kinetic energy: Energy of motion.

• Units: calorie (cal), kilocalorie (kcal), joule (J), kilojoule (kJ).

• $1\ \text{cal} = 4.184\ \text{J}$

3.5 Specific Heat and Heat Calculations

• Specific heat: Energy required to raise the temperature of 1 g of a substance by 1°C.

• Formula: $q = m \times c \times \Delta T$ where $q$ = heat (J or cal), $m$ = mass (g), $c$ = specific heat, $\Delta T$ = temperature change (°C)

3.6 Food Energy Calculations

• Nutritional Calorie (Cal): 1 Cal = 1000 cal = 1 kcal.

• Food energy is the sum of kcal or kJ from carbohydrates, fats, and proteins.

3.7 Changes of State and Energy

• Melting: Solid to liquid; requires heat of fusion.

• Freezing: Liquid to solid; releases heat of fusion.

• Evaporation/Boiling: Liquid to gas; requires heat of vaporization.

• Condensation: Gas to liquid; releases heat of vaporization.

• Sublimation: Solid to gas directly.

• Heating/Cooling Curve: Graph showing temperature changes and plateaus during state changes.

• Total heat change: Sum of energy for temperature changes and state changes.

Ch 4: Atoms & Elements

4.1 Element Names and Symbols

• Elements: Fundamental substances of matter.

• Chemical symbols: One- or two-letter abbreviations (e.g., H for hydrogen, Na for sodium).

4.2 The Periodic Table: Groups, Periods, and Classifications

• Periodic table: Arranged by increasing atomic number.

• Period: Horizontal row.

• Group: Vertical column; elements in a group have similar properties.

• Special groups: Alkali metals (1A/1), alkaline earth metals (2A/2), halogens (7A/17), noble gases (8A/18).

• Metals: Left of zigzag line; Nonmetals: Right of zigzag; Metalloids: Along zigzag (except Al, Og).

4.3 Subatomic Particles and Atomic Structure

• Atom: Smallest unit of an element.

• Proton: Positive charge (+), in nucleus.

• Neutron: No charge, in nucleus.

• Electron: Negative charge (–), outside nucleus.

4.4 Atomic Number, Mass Number, and Isotopes

• Atomic number (Z): Number of protons.

• Mass number (A): Protons + neutrons.

• In a neutral atom, protons = electrons.

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

• Atomic mass: Weighted average of all isotopes.

4.6 Electron Arrangement

• Electrons occupy energy levels (shells) around the nucleus.

• First level: 2 electrons; second: 8; third: up to 18.

• Electron arrangement is written by filling lower energy levels first.

4.7 Periodic Trends

• Valence electrons: Outermost electrons; determine chemical properties.

• Elements in a group have similar valence electron arrangements.

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

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

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

Ch 5: Nuclear Chemistry

5.1 Types of Radiation

• Radioactive isotopes: Unstable nuclei that decay, emitting radiation.

• Alpha (α): Helium nucleus; low penetration.

• Beta (β): Electron; moderate penetration.

• Positron (β+): Positive electron; moderate penetration.

• Gamma (γ): High-energy photon; high penetration.

• Shielding is required to protect from radiation.

5.2 Nuclear Equations

• Balanced nuclear equations show changes in mass and atomic numbers during decay.

• Symbols indicate isotopes and emitted particles.

• Artificial radioisotopes are produced by bombarding stable isotopes with particles.

• Example: $^{14}_{6}\mathrm{C} \rightarrow ^{14}_{7}\mathrm{N} + ^{0}_{-1}\beta$ (beta decay)

5.3 Detection and Measurement of Radiation

• Geiger counter: Detects radiation via ionization of gas, producing an electrical current.

• Activity: Measured in curie (Ci) or becquerel (Bq); number of disintegrations per second.

• Absorbed dose: Measured in rad or gray (Gy).

• Biological damage: Measured in rem or sievert (Sv).

5.4 Half-Life Calculations

• Half-life: Time for half of a radioactive sample to decay.

• Short half-lives: Medical isotopes (e.g., Tc-99, I-131).

• Long half-lives: Environmental isotopes (e.g., C-14, Ra-226, U-238).

• Calculation: After $n$ half-lives, remaining amount = $(\frac{1}{2})^n \times$ original amount.

5.5 Medical Uses of Radioisotopes

• Radioisotopes target specific body sites for diagnosis or treatment.

• Radiation detection helps evaluate organ function, disease, or tumors.

• High doses can treat or destroy tumors.

5.6 Nuclear Fission and Fusion

• Fission: Large nucleus splits into smaller nuclei, releasing energy and radiation.

• Fusion: Small nuclei combine to form a larger nucleus, releasing even more energy.

• Example: Fission: $^{235}_{92}\mathrm{U} + ^{1}_{0}\mathrm{n} \rightarrow ^{141}_{56}\mathrm{Ba} + ^{92}_{36}\mathrm{Kr} + 3^{1}_{0}\mathrm{n}$

• Fusion: $^{2}_{1}\mathrm{H} + ^{3}_{1}\mathrm{H} \rightarrow ^{4}_{2}\mathrm{He} + ^{1}_{0}\mathrm{n}$