CHE103 Exam 1 Study Guide: Chapters 1-5 (Chemistry in Our Lives, Measurements, Matter & Energy, Atoms & Elements, Nuclear Chemistry)

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Chapter 1: Chemistry in Our Lives

Matter and Chemicals

Chemistry is the study of matter, which is anything that has mass and occupies space. Matter can be classified based on whether it contains chemicals, which are substances with a definite composition.

Matter: Anything that has mass and takes up space.
Chemical: A substance with a definite composition, such as water (H2O), sodium chloride (NaCl).
Example: Air, water, and table salt are all examples of matter containing chemicals.

The Scientific Method

The scientific method is a systematic approach to investigation in science.

Observation: Gathering information using the senses.
Hypothesis: A tentative explanation for an observation.
Experiment: A procedure to test the hypothesis.
Conclusion: A decision based on the results of the experiment.
Example: Observing that plants grow toward light, hypothesizing that light affects growth, testing with different light sources, and concluding based on results.

Numerical Skills in Chemistry

Place Value: Understanding the value of digits in a number (hundreds, tens, ones, tenths, hundredths, etc.).
Percentage Calculation: $\text{Percentage} = \frac{\text{Part}}{\text{Total}} \times 100\%$
Graph Interpretation: Reading and extracting information from graphical data.
Basic Math Operations: Addition (+), subtraction (−), multiplication (×), division (÷).
Scientific Notation: Expressing numbers as a product of a coefficient and a power of ten, e.g., $3.2 \times 10^4$.

Chapter 2: Chemistry & Measurement

Exact vs. Measured Numbers

Exact Number: A value known with complete certainty (e.g., 12 eggs in a dozen).
Measured Number: A value obtained by measurement, subject to uncertainty.

Significant Figures (SFs)

Significant figures reflect the precision of a measured value.

Reporting SFs: Record all certain digits plus one estimated digit.
Rules for Calculations:
- Addition/Subtraction: Result has the same number of decimal places as the measurement with the fewest decimal places.
- Multiplication/Division: Result has the same number of significant figures as the measurement with the fewest SFs.

Unit Conversions

Metric Prefixes: kilo (k, $10^3$), deci (d, $10^{-1}$), centi (c, $10^{-2}$), milli (m, $10^{-3}$), micro ($\mu$, $10^{-6}$).
Common Equalities:
- 1 kg = 1000 g
- 1 g = 1000 mg
- 1 mg = 1000 mcg
- 1 L = 1000 mL
- 1 mL = 1 cm3
- 1 km = 1000 m
- 1 m = 100 cm
- 1 cm = 10 mm
- 1 m = 1000 mm
- 1 hour = 60 min
- 1 min = 60 s
Conversion Factors: Ratios derived from equalities, used to convert units.
Example: To convert 5.0 cm to mm: $5.0\,\text{cm} \times \frac{10\,\text{mm}}{1\,\text{cm}} = 50\,\text{mm}$

Density

Definition: Density is mass per unit volume.
Formula: $\text{Density} = \frac{\text{Mass (g)}}{\text{Volume (mL)}}$
Example: If mass = 10 g and volume = 2 mL, density = $\frac{10}{2} = 5\,\text{g/mL}$

Chapter 3: Matter & Energy

Classification of Matter

Pure Substance: Has a fixed composition; can be an element or a compound.
Mixture: Contains two or more substances physically combined; can be homogeneous (uniform) or heterogeneous (non-uniform).
Example: Salt water is a homogeneous mixture; sand and iron filings is a heterogeneous mixture.

States of Matter

Solid: Definite shape and volume; particles are closely packed.
Liquid: Definite volume, no definite shape; particles are less tightly packed.
Gas: No definite shape or volume; particles are far apart.

Chemical vs. Physical Changes and Properties

Physical Change: Change in state or appearance without changing composition (e.g., melting ice).
Chemical Change: Change that produces new substances (e.g., rusting iron).
Physical Property: Observable without changing substance (e.g., color, melting point).
Chemical Property: Describes ability to change into new substances (e.g., flammability).

Temperature Conversions

°C to °F: $T_F = 1.8T_C + 32$
°F to °C: $T_C = \frac{T_F - 32}{1.8}$
°C to K: $T_K = T_C + 273$
K to °C: $T_C = T_K - 273$

Energy and Its Units

Potential Energy: Stored energy due to position.
Kinetic Energy: Energy of motion.
Unit Conversions:
- 1 cal = 4.184 J
- 1 kJ = 1000 J
- 1 kcal = 1000 cal

Energy in Food

Energy Value (kJ/g): $\text{Energy Value} = \frac{\text{Energy (kJ)}}{\text{mass (g)}}$
Total Energy: $\text{Energy (kJ)} = \text{Energy Value (kJ/g)} \times \text{mass (g)}$
Example: If a food contains 10 g of carbohydrate (17 kJ/g), energy = $10 \times 17 = 170$ kJ

Specific Heat

Definition: Amount of heat required to raise the temperature of 1 g of a substance by 1°C.
Formula: $\text{Specific Heat} = \frac{\text{Heat (J)}}{\text{mass (g)} \times \Delta T (°C)}$
Rearranged:
- To solve for heat: $\text{Heat (J)} = \text{Specific Heat} \times \text{mass (g)} \times \Delta T (°C)$
- To solve for change in temperature: $\Delta T (°C) = \frac{\text{Heat (J)}}{\text{Specific Heat} \times \text{mass (g)}}$

Changes of State

Melting: Solid to liquid
Freezing: Liquid to solid
Boiling: Liquid to gas
Condensation: Gas to liquid

Heating and Cooling Curves

Heating Curve: Shows temperature change as heat is added; plateaus indicate phase changes (melting point, boiling point).
Cooling Curve: Shows temperature change as heat is removed.

Chapter 4: Atoms & Elements

Elements and Symbols

Element: Pure substance consisting of one type of atom.
Element Symbols: One- or two-letter abbreviations (e.g., H for hydrogen, Au for gold).
Common Elements: Hydrogen (H), Helium (He), Lithium (Li), Beryllium (Be), ... Krypton (Kr), Gold (Au), Silver (Ag), Mercury (Hg), Lead (Pb), Tin (Sn), Platinum (Pt).

The Periodic Table

Groups: Vertical columns; Group 1 (1A) - Alkali metals, Group 2 (2A) - Alkaline earth metals, Group 17 (7A) - Halogens, Group 18 (8A) - Noble gases.
Periods: Horizontal rows.
Classification: Metals, nonmetals, metalloids.

Atomic Structure

Proton: Positive charge (+1), located in nucleus.
Neutron: No charge (0), located in nucleus.
Electron: Negative charge (−1), located outside nucleus.

Atomic Number, Mass Number, Atomic Weight

Atomic Number (Z): Number of protons in nucleus.
Mass Number (A): Number of protons plus neutrons.
Atomic Weight: Weighted average mass of all isotopes.
Atomic Symbol: Notation showing element, atomic number, and mass number (e.g., $^{23}_{11}\text{Na}$).

Isotopes

Definition: Atoms of the same element with different numbers of neutrons.
Example: Carbon-12 and Carbon-14 are isotopes of carbon.

Calculating Subatomic Particles

Protons: Equal to atomic number.
Neutrons: Mass number minus atomic number.
Electrons: Equal to protons in a neutral atom.

Electron Arrangement and Valence Electrons

Electron Arrangement: Distribution of electrons in shells around nucleus (first 20 elements).
Valence Electrons: Electrons in the outermost shell; determine chemical reactivity.
Lewis Symbol: Element symbol with dots representing valence electrons.

Periodic Trends

Atomic Size: Increases down a group, decreases across a period.
Metallic Character: Increases down a group, decreases across a period.
Ionization Energy: Energy required to remove an electron; decreases down a group, increases across a period.

Chapter 5: Nuclear Chemistry

Types of Radiation

Alpha Particle (α): $^4_2\text{He}$; mass number 4, atomic number 2.
Beta Particle (β): $^0_{-1}\text{e}$; mass number 0, atomic number −1.
Positron (β+): $^0_{+1}\text{e}$; mass number 0, atomic number +1.
Gamma Ray (γ): $^0_0\gamma$; no mass or charge.

Biological Effects and Protection

Radiation Exposure: Can damage living tissue; effects depend on type and amount.
Protection: Use shielding (lead, concrete), minimize exposure time, increase distance from source.

Radioactive Decay Equations

Balanced Nuclear Equation: Shows parent and daughter nuclides, mass and atomic numbers balanced.
Example: $^{14}_6\text{C} \rightarrow ^{14}_7\text{N} + ^0_{-1}\text{e}$ (beta decay)

Half-Life Calculations

Half-Life (t1/2): Time required for half of a radioactive sample to decay.
Amount Remaining: $\text{Remaining} = \text{Initial} \times \left(\frac{1}{2}\right)^n$ where n = number of half-lives.
Number of Half-Lives: $n = \frac{\text{Total Time}}{\text{Half-Life}}$
Example: If 100 g of a radioisotope with a half-life of 3 years, after 6 years: $100 \times (1/2)^2 = 25$ g remain.

Summary of Key Equations and Conversion Factors

Equality	Conversion Factor
1 ft = 12 in.	1 ft / 12 in. or 12 in. / 1 ft
1 in. = 2.54 cm	1 in. / 2.54 cm or 2.54 cm / 1 in.
1 m = 39.4 in.	1 m / 39.4 in. or 39.4 in. / 1 m
1 kg = 2.20 lb	1 kg / 2.20 lb or 2.20 lb / 1 kg
Density	$\text{Density} = \frac{\text{Mass (g)}}{\text{Volume (mL)}}$
Temperature (°F)	$T_F = 1.8T_C + 32$
Temperature (°C)	$T_C = \frac{T_F - 32}{1.8}$
1 cal = 4.184 J	1 cal / 4.184 J or 4.184 J / 1 cal
Specific Heat	$\text{Specific Heat} = \frac{\text{Heat (J)}}{\text{mass (g)} \times \Delta T (°C)}$
1 kg = 1000 g	1 kg / 1000 g or 1000 g / 1 kg
1 g = 1000 mg	1 g / 1000 mg or 1000 mg / 1 g
1 mg = 1000 mcg	1 mg / 1000 mcg or 1000 mcg / 1 mg
1 L = 1000 mL	1 L / 1000 mL or 1000 mL / 1 L
1 mL = 1 cm3	1 mL / 1 cm3
1 km = 1000 m	1 km / 1000 m or 1000 m / 1 km
1 m = 100 cm	1 m / 100 cm or 100 cm / 1 m
1 cm = 10 mm	1 cm / 10 mm or 10 mm / 1 cm
1 m = 1000 mm	1 m / 1000 mm or 1000 mm / 1 m
1 hour = 60 min	1 hour / 60 min or 60 min / 1 hour
1 min = 60 s	1 min / 60 s or 60 s / 1 min
Temperature (K)	$T_K = T_C + 273$
Temperature (°C)	$T_C = T_K - 273$
1 kJ = 1000 J	1 kJ / 1000 J or 1000 J / 1 kJ
1 kcal = 1000 cal	1 kcal / 1000 cal or 1000 cal / 1 kcal
Energy Value	$\text{Energy Value (kJ/g)} = \frac{\text{Energy (kJ)}}{\text{mass (g)}}$
Energy	$\text{Energy (kJ)} = \text{Energy Value (kJ/g)} \times \text{mass (g)}$
Percentage	$\text{Percentage} = \frac{\text{Part}}{\text{Total}} \times 100\%$