Introductory Chemistry: Measurement, Matter, and Atomic Structure Study Guide

Chapter 2: Measurement and Problem Solving

Key Vocabulary and Concepts

• Exact Number: A value known with complete certainty, often from counting or defined quantities (e.g., 1 dozen = 12).

• Measurement: A quantitative observation that includes both a number and a unit (e.g., 25.0 mL).

• Precision: The degree to which repeated measurements under unchanged conditions show the same results.

• Accuracy: How close a measured value is to the true or accepted value.

• Significant Figures (Sig Figs): The digits in a measurement that are known with certainty plus one estimated digit.

• SI Units: The International System of Units, the standard units used in science (e.g., meter, kilogram, second).

• Mass: The amount of matter in an object, measured in kilograms (kg).

• Volume: The amount of space an object occupies, measured in liters (L) or cubic centimeters (cm3).

• Density: The mass per unit volume of a substance, calculated as $\text{Density} = \frac{\text{Mass}}{\text{Volume}}$.

• Unit Analysis (Dimensional Analysis): A method for converting between units using conversion factors.

• Conversion Factor: A ratio used to express a quantity in different units.

• Uncertainty: The degree of doubt in a measurement, often indicated by the last significant figure.

Applying Accuracy and Precision

• Accuracy refers to how close a measurement is to the true value; precision refers to how close repeated measurements are to each other.

• Example: If you weigh a 10.00 g standard mass and obtain values of 9.98 g, 9.99 g, and 10.01 g, your measurements are both accurate and precise.

Scientific Notation

• Used to express very large or very small numbers in the form $a \times 10^n$.

• Example: 0.00045 = $4.5 \times 10^{-4}$

Significant Figures in Calculations

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

• Multiplication/Division: The result should have the same number of significant figures as the measurement with the fewest significant figures.

• Ambiguous Zeros: Zeros between nonzero digits are significant; leading zeros are not; trailing zeros are significant only if there is a decimal point.

Rounding Calculated Values

• Round the final answer to the correct number of significant figures based on the rules above.

Unit Conversions

• Use conversion factors to change from one unit to another (e.g., inches to centimeters).

• Metric prefixes (kilo-, centi-, milli-, etc.) are used to indicate multiples or fractions of units.

• Area and volume conversions often require squaring or cubing the conversion factor.

Density Calculations

• Density is calculated as $\text{Density} = \frac{\text{Mass}}{\text{Volume}}$.

• Density can be used as a conversion factor between mass and volume.

SI Units for Common Quantities

• Mass: kilogram (kg)

• Volume: liter (L)

• Time: second (s)

• Temperature: kelvin (K)

Chapter 3: Matter and Energy

Classification of Matter

• Matter: Anything that has mass and occupies space.

• Pure Substance: Matter with a fixed composition (element or compound).

• Mixture: A combination of two or more substances that are not chemically combined.

• Homogeneous Mixture (Solution): Uniform composition throughout (e.g., saltwater).

• Heterogeneous Mixture: Non-uniform composition (e.g., salad, sand and iron filings).

• Element: A pure substance made of only one kind of atom.

• Compound: A pure substance composed of two or more elements chemically combined.

Physical and Chemical Properties and Changes

• Physical Properties: Characteristics that can be observed without changing the substance (e.g., color, melting point).

• Physical Change: A change that does not alter the chemical composition (e.g., melting, boiling).

• Chemical Properties: Characteristics that describe a substance's ability to undergo a chemical change (e.g., flammability).

• Chemical Change: A change that produces one or more new substances (e.g., rusting, burning).

• Visible signs of chemical change: color change, gas production, precipitate formation, temperature change.

Law of Conservation of Mass

• Mass is neither created nor destroyed in a chemical reaction; the total mass of reactants equals the total mass of products.

Energy and Its Forms

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

• Kinetic Energy: Energy of motion.

• Potential Energy: Stored energy due to position or composition.

• Other forms: electrical, thermal, chemical energy.

Energy Units and Conversions

• Common units: joule (J), calorie (cal), kilowatt-hour (kWh).

• 1 calorie (cal) = 4.184 joules (J)

• 1 Calorie (Cal) = 1000 calories (cal)

• 1 kilowatt-hour (kWh) = $3.60 \times 10^6$ joules (J)

Exothermic vs. Endothermic Processes

• Exothermic: Processes that release energy (heat flows out of the system).

• Endothermic: Processes that absorb energy (heat flows into the system).

Temperature and Particle Motion

• Temperature is a measure of the average kinetic energy of particles in a substance.

• As temperature increases, particle motion increases.

Temperature Scales and Conversions

• Celsius (°C), Fahrenheit (°F), Kelvin (K)

• Conversions:

  • $T_{\text{K}} = T_{\text{°C}} + 273.15$

  • $T_{\text{°C}} = \frac{5}{9}(T_{\text{°F}} - 32)$

  • $T_{\text{°F}} = \frac{9}{5}T_{\text{°C}} + 32$

Specific Heat Capacity

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

• Equation: $q = m \times C \times \Delta T$

• Where $q$ = heat (J), $m$ = mass (g), $C$ = specific heat capacity (J/g·°C), $\Delta T$ = change in temperature (°C).

Chapter 4: Atoms and Elements

Key Vocabulary

• Electron: Negatively charged subatomic particle found outside the nucleus.

• Proton: Positively charged subatomic particle found in the nucleus.

• Neutron: Neutral subatomic particle found in the nucleus.

• Nucleus: Dense center of the atom containing protons and neutrons.

• Atomic Number (Z): Number of protons in the nucleus; defines the element.

• Mass Number (A): Total number of protons and neutrons in the nucleus.

• Atomic Mass: Weighted average mass of the atoms in a naturally occurring element.

• Atomic Mass Unit (amu): Standard unit for atomic mass; 1 amu = 1/12 the mass of a carbon-12 atom.

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

• Ions: Atoms or molecules with a net electric charge due to loss or gain of electrons.

• Cation: Positively charged ion (loss of electrons).

• Anion: Negatively charged ion (gain of electrons).

• Periodicity: The recurring trends in properties of elements across periods of the periodic table.

• Group: Vertical column in the periodic table.

• Period: Horizontal row in the periodic table.

• Transition Metals: Elements in groups 3-12.

• Lanthanide and Actinide Series: Two rows below the main body of the periodic table.

• Noble Gases: Group 18 elements, inert gases.

• Metals, Nonmetals, Metalloids: Classification based on properties and location in the periodic table.

• Alkali Metals: Group 1 elements.

• Alkaline Earth Metals: Group 2 elements.

• Chalcogens: Group 16 elements.

• Halogens: Group 17 elements.

Law of Conservation of Mass in Chemical Reactions

• The total mass of reactants equals the total mass of products in a chemical reaction.

Isotopes and Atomic Structure

• Isotopes are identified by their mass number (e.g., carbon-12, carbon-14).

• Number of protons = atomic number; number of neutrons = mass number - atomic number; number of electrons = number of protons (for neutral atoms).

Ions and Their Charges

• Main group elements tend to form ions with predictable charges (e.g., Group 1 forms +1, Group 17 forms -1).

• Cations are formed by losing electrons; anions are formed by gaining electrons.

Classification of Elements

• Representative (main group) elements: Groups 1, 2, and 13-18.

• Transition elements: Groups 3-12.

• Metals: Left and center of the periodic table; good conductors, malleable, ductile.

• Nonmetals: Right side; poor conductors, brittle.

• Metalloids: Properties intermediate between metals and nonmetals; found along the stair-step line.

Calculating Atomic Mass from Isotopic Abundances

• Atomic mass is calculated as the weighted average of the masses of all naturally occurring isotopes.

• Equation: $\text{Atomic Mass} = \sum (\text{Fractional Abundance} \times \text{Isotopic Mass})$

Essential Memorization

• Element names and symbols for elements #1-56 and 72-86.

• Metric prefixes, symbols, and exponents.

Tables for Reference

• Common unit conversions for length, mass, and volume.

• Densities of common substances.

• Energy conversion factors.

• Specific heat capacities of common substances.

Additional info: Students should be able to interpret and analyze data from tables and graphs, as well as apply these concepts to solve quantitative problems in chemistry.