Chem 410 Exam 1 Study Guide: Chapters 1–4 (Chemistry in Our Lives, Measurements, Matter & Energy, Atoms & Elements)

Chapter 1: Chemistry in Our Lives

Scientific Method

The scientific method is a systematic approach used by scientists to investigate phenomena, acquire new knowledge, or correct and integrate previous knowledge.

• Observation: Gathering information through senses or instruments.

• Hypothesis: A tentative explanation for observations.

• Experiment: Testing the hypothesis under controlled conditions.

• Conclusion: Analyzing results to accept, reject, or modify the hypothesis.

• Theory: A well-substantiated explanation based on repeated experiments.

What is a Chemical?

A chemical is any substance that has a definite composition. Examples include water (H2O), table salt (NaCl), and glucose (C6H12O6).

• All matter is made of chemicals.

• Chemicals can be natural or synthetic.

Solving Equations & Place Value

Understanding the place value of numbers is essential for solving equations and performing calculations in chemistry.

• Units: Ones, tens, hundreds, thousands, etc.

• Example: In 345, the digit 3 is in the hundreds place.

Chapter 2: Chemistry and Measurements

Exact Numbers vs. Measured Numbers

Exact numbers are values known with complete certainty (e.g., counting objects), while measured numbers are obtained using instruments and have some degree of uncertainty.

• Exact: 12 eggs in a dozen.

• Measured: 5.67 cm (using a ruler).

Unit Conversions

Unit conversions are essential for expressing measurements in different units. The Metric/SI system uses prefixes to indicate multiples or fractions of units.

• Mass: grams (g), kilograms (kg), milligrams (mg)

• Length: meters (m), centimeters (cm), millimeters (mm)

• Temperature: Celsius (°C), Kelvin (K)

• Common prefixes: kilo- (103), centi- (10-2), milli- (10-3), micro- (10-6)

• Unit canceling: Use conversion factors to cancel units and solve problems.

Significant Figures (Sig Figs)

Significant figures indicate the precision of a measured value.

• Rules: All nonzero digits are significant; zeros between nonzero digits are significant; leading zeros are not significant; trailing zeros are significant if there is a decimal point.

• Scientific notation: Expresses numbers as a product of a coefficient and a power of ten (e.g., $3.2 \times 10^4$).

• Multiplication/Division: The answer should have the same number of sig figs as the measurement with the fewest sig figs.

• Addition/Subtraction: The answer should have the same number of decimal places as the measurement with the fewest decimal places.

Density, Mass, and Volume

Density is the mass of a substance per unit volume.

• Formula: $\text{Density} = \frac{\text{Mass}}{\text{Volume}}$

• Units: g/mL, g/cm3

• Example: If mass = 10 g and volume = 2 mL, density = 5 g/mL.

Chapter 3: Matter and Energy

Classification of Matter

Matter can be classified based on its composition.

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

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

• Mixture: Physical combination of two or more substances.

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

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

States and Properties of Matter

Matter exists in different states and can undergo physical or chemical changes.

• States: Solid, liquid, gas.

• Physical change: Change in state or appearance without altering composition (e.g., melting ice).

• Chemical change: Change that produces new substances (e.g., burning wood).

Energy Content of Food

Food provides energy through carbohydrates, fats, and proteins.

• Carbohydrates: 4 kcal/g

• Fats: 9 kcal/g

• Proteins: 4 kcal/g

• Formula: $\text{Total Energy} = (\text{g carbs} \times 4) + (\text{g fat} \times 9) + (\text{g protein} \times 4)$

• Example: 10 g carbs, 5 g fat, 8 g protein: $(10 \times 4) + (5 \times 9) + (8 \times 4) = 40 + 45 + 32 = 117$ kcal

Specific Heat

Specific heat is the amount of heat required to raise the temperature of 1 gram of a substance by 1°C.

• Formula: $q = m \times c \times \Delta T$

• Where: q = heat (J), m = mass (g), c = specific heat (J/g°C), ΔT = change in temperature (°C)

• Units: J, kJ, cal, kcal, Cal

• Example: 50 g water, c = 4.18 J/g°C, ΔT = 10°C: $q = 50 \times 4.18 \times 10 = 2090$ J

Heating Curve

A heating curve shows the temperature change of a substance as heat is added.

• Heat of Fusion (Melting): Energy required to change a solid to a liquid.

• Heat of Vaporization (Boiling): Energy required to change a liquid to a gas.

• Specific heat: Energy required to change temperature within a phase.

Chapter 4: Atoms and Elements

Periodic Table

The periodic table organizes elements by increasing atomic number and groups elements with similar properties.

• Groups: Vertical columns; elements in a group have similar properties.

• Periods: Horizontal rows.

• Types: Metals, nonmetals, metalloids, noble gases.

• Example: Group 1: Alkali metals; Group 18: Noble gases.

Identifying Elements

Elements are identified by their symbol, atomic number, and the number of protons, neutrons, and electrons.

• Atomic number: Number of protons.

• Mass number: Protons + neutrons.

• Example: Carbon (C): Atomic number = 6, Mass number = 12.

Structure of the Atom

An atom consists of subatomic particles: protons, neutrons, and electrons.

• Proton: Positive charge (+1), mass ≈ 1 amu.

• Neutron: Neutral charge (0), mass ≈ 1 amu.

• Electron: Negative charge (-1), mass ≈ 0.0005 amu.

Isotopes

Isotopes are atoms of the same element with different numbers of neutrons.

• Atomic symbol: $^{A}_{Z} \text{X}$ where A = mass number, Z = atomic number, X = element symbol.

• Example: $^{14}_{6} \text{C}$ is an isotope of carbon.

• Calculating atomic mass: $\text{Atomic mass} = \sum (\text{isotope mass} \times \text{abundance})$

Atomic Structure: Energy Levels and Orbitals

Electrons occupy energy levels and sublevels (s, p, d, f) with specific shapes and orientations.

• Principal quantum number (n): Indicates energy level.

• Sublevels: s (sphere), p (dumbbell), d, f (complex shapes).

• Electrons per sublevel: s = 2, p = 6, d = 10, f = 14.

• Electron spin: Each orbital can hold two electrons with opposite spins.

Electron Configurations

Electron configuration describes the arrangement of electrons in an atom.

• Orbital diagrams: Visual representation of electron placement.

• Full configuration: Lists all occupied sublevels (e.g., 1s2 2s2 2p6).

• Abbreviated configuration: Uses noble gas core (e.g., [Ne] 3s2).

• Outer shell configuration: Shows valence electrons.

Valence Electrons

Valence electrons are electrons in the outermost energy level and determine chemical reactivity.

• Group number indicates number of valence electrons (e.g., Group 1 = 1 valence electron).

• Valence electron configuration is important for bonding.

Electron-Dot Symbols

Electron-dot symbols (Lewis structures) represent valence electrons as dots around the element symbol.

• 1–4 valence electrons: single dots around the symbol.

• 5–8 valence electrons: some dots paired.

• Example: Oxygen (6 valence electrons): O with two pairs and two single dots.

Periodic Trends

Periodic trends describe how properties change across periods and groups.

• Atomic size: Increases down a group, decreases across a period (left to right).

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

• Explanation: Atomic size increases due to more energy levels; ionization energy increases due to greater nuclear charge.

Property	Trend Down Group	Trend Across Period
Atomic Size	Increases	Decreases
Ionization Energy	Decreases	Increases