Introductory Chemistry Exam 1 Study Guide

The Chemical World

Where Are Chemicals Found?

Chemicals are present everywhere in our daily lives, including in the air we breathe, the food we eat, and the products we use. Chemistry studies the composition, structure, and changes of matter.

• Examples: Water (H2O), table salt (NaCl), and oxygen gas (O2) are all chemicals.

Measurement and Problem Solving

Scientific Notation

Scientific notation expresses very large or very small numbers in the form a × 10n.

• Example: 0.00056 = 5.6 × 10-4

Unit Conversions

Unit conversions use conversion factors to change from one unit to another.

• Example: To convert 5.0 cm to meters: $5.0\ \mathrm{cm} \times \frac{1\ \mathrm{m}}{100\ \mathrm{cm}} = 0.050\ \mathrm{m}$

Calculations with Correct Significant Figures

Significant figures reflect the precision of a measurement. The result of a calculation should be reported with the correct number of significant figures.

• Rules: For multiplication/division, the answer has the same number of significant figures as the measurement with the fewest significant figures. For addition/subtraction, the answer has the same number of decimal places as the measurement with the fewest decimal places.

Using Measuring Devices

Measurements should be recorded to the correct number of significant figures, including all certain digits plus one estimated digit.

Interpreting Data from a Line Graph

Line graphs display the relationship between two variables. The slope and shape of the line can indicate trends, such as direct or inverse relationships.

Matter and Energy

Physical and Chemical Changes

Physical changes do not alter the chemical composition of a substance (e.g., melting, boiling). Chemical changes result in the formation of new substances (e.g., rusting, burning).

• Example: Ice melting is a physical change; iron rusting is a chemical change.

Properties of Solids, Liquids, and Gases

• Solids: Definite shape and volume; particles are closely packed.

• Liquids: Definite volume, indefinite shape; particles can move past each other.

• Gases: Indefinite shape and volume; particles are far apart and move freely.

Properties of Compounds, Elements, and Mixtures

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

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

• Mixture: Physical blend of two or more substances. Homogeneous mixtures are uniform throughout (e.g., saltwater); heterogeneous mixtures are not uniform (e.g., salad).

Properties of Water

• Boiling Point: 100°C (373 K) at 1 atm

• Melting Point: 0°C (273 K) at 1 atm

Atoms and Elements

Atomic Structure

Atoms consist of protons, neutrons, and electrons.

• Protons: Positively charged, found in the nucleus

• Neutrons: Neutral, found in the nucleus

• Electrons: Negatively charged, found outside the nucleus

Atomic Number, Mass Number, and Atomic Mass

• Atomic Number (Z): Number of protons in the nucleus

• Mass Number (A): Total number of protons and neutrons

• Atomic Mass: Weighted average mass of all isotopes of an element

Isotopes

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

• Example: Carbon-12 and Carbon-14 are isotopes of carbon.

Calculating Atomic Mass from Isotopic Distribution

The atomic mass is calculated using the relative abundance and mass of each isotope:

$\text{Atomic mass} = \sum (\text{fractional abundance} \times \text{isotope mass})$

Dalton’s Atomic Theory

• All matter is composed of atoms.

• Atoms of the same element are identical.

• Atoms cannot be created or destroyed in chemical reactions.

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

Conservation of Mass

Mass is neither created nor destroyed in a chemical reaction.

Elements That Occur as Diatomic Molecules or Atoms

• Diatomic elements: H2, N2, O2, F2, Cl2, Br2, I2

• Monatomic elements: Noble gases (He, Ne, Ar, Kr, Xe, Rn)

Chemical Composition

Cations and Anions

• Cation: Positively charged ion (loss of electrons)

• Anion: Negatively charged ion (gain of electrons)

Expected Charges of Ions

• Group 1 (Alkali metals): +1

• Group 2 (Alkaline earth metals): +2

• Halogens (Group 17): -1

Symbols for Elements and Writing Charges on Ions

• Example: Sodium ion: Na+; Chloride ion: Cl-

Electrons in Atoms and the Periodic Table

Identifying Groups of Elements

Group	Location	Properties
Alkali metals	Group 1	Very reactive, form +1 ions
Alkaline earth metals	Group 2	Reactive, form +2 ions
Transition metals	Groups 3-12	Variable charges, often form colored compounds
Main-group elements	Groups 1, 2, 13-18	Wide range of properties
Halogens	Group 17	Very reactive nonmetals, form -1 ions
Noble gases	Group 18	Inert, rarely form compounds

Properties of Metals, Nonmetals, and Metalloids

Type	Properties
Metals	Shiny, good conductors, malleable, ductile
Nonmetals	Dull, poor conductors, brittle
Metalloids	Intermediate properties, semiconductors

Chemical Reactions and Energy

Energy Diagrams: Exothermic vs. Endothermic Reactions

• Exothermic: Releases energy; products have lower energy than reactants.

• Endothermic: Absorbs energy; products have higher energy than reactants.

Specific Heat Capacity Calculations

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

Equation:

$q = m \times C \times \Delta T$

• q: Heat (Joules)

• m: Mass (grams)

• C: Specific heat capacity (J/g·°C)

• ΔT: Change in temperature (°C)

Temperature Conversions

• Celsius to Kelvin: $K = ^\circ C + 273.15$

• Kelvin to Celsius: $^\circ C = K - 273.15$

Scientific Method and Laboratory Safety

Laws, Theories, Observations, Hypotheses

• Observation: Gathering information using senses or instruments.

• Hypothesis: Testable explanation for an observation.

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

• Law: Statement that describes a consistent relationship in nature.

Basic Laboratory Safety Hazards

• Wear appropriate protective equipment (goggles, gloves, lab coat).

Know the location of safety equipment (eyewash, fire extinguisher).

• Never eat or drink in the laboratory.

• Dispose of chemicals properly.