CHEM 1201: Basic Chemistry – Structure, Properties, and Measurement

Energy Terminology and Thermodynamics

First Law of Thermodynamics and Energy Conservation

The study of energy in chemistry is governed by the First Law of Thermodynamics, which states that energy in the universe is conserved. This means that energy can neither be created nor destroyed, only transferred or transformed.

• System: The part of the universe under study (e.g., a chemical reaction or phase change).

• Surroundings: Everything outside the system.

• Universe: The sum of the system and its surroundings.

Diagram: The system is typically represented as a region within a larger area representing the surroundings.

Exothermic vs. Endothermic Processes

Chemical and physical processes can be classified based on the direction of heat (energy) flow between the system and surroundings.

• Exothermic Process:

  • The system loses heat to the surroundings.

  • The energy of the system decreases, while the energy of the surroundings increases.

  • Heat is assigned a negative sign ($q < 0$).

  • Example: Combustion of natural gas in a stove.

• Endothermic Process:

  • The system gains heat from the surroundings.

  • The energy of the system increases, while the energy of the surroundings decreases.

  • Heat is assigned a positive sign ($q > 0$).

  • Example: Melting of ice.

Conceptual Example

Is the burning of natural gas in a stove exothermic or endothermic? What is the sign of the energy change?

• Answer: Exothermic; negative sign.

Dimensional Analysis and Problem Solving

Dimensional Analysis: Strategy for Solving Problems

Dimensional analysis is a systematic approach to problem solving that uses units as a guide. It is essential for converting between different units of measurement in chemistry.

• Conversion Factor: A fractional quantity expressing the relationship between two units. The unit to be converted from is placed in the denominator, and the unit to be converted to is placed in the numerator.

• Equivalence Statement: A statement that shows two quantities are equal (e.g., 1 inch = 2.54 cm).

Conversion factors can be written as:

• $\frac{1\ in}{2.54\ cm}$

• $\frac{2.54\ cm}{1\ in}$

Arrange conversion factors so that the starting unit cancels, and the desired unit remains.

• Example: To convert 2.00 ft to inches: $2.00\ \text{ft} \times \frac{12\ \text{in}}{1\ \text{ft}} = 24.0\ \text{inches}$

Classification of Matter

States of Matter

Matter can be classified by its physical state: solid, liquid, or gas. The state depends on the arrangement and movement of particles.

• Solid: Atoms or molecules are closely packed in fixed positions. Solids have a fixed volume and shape. Examples: Ice, aluminum, diamond.

• Liquid: Atoms or molecules are close but can move past each other. Liquids have a fixed volume but take the shape of their container. Examples: Water, alcohol, gasoline.

• Gas: Atoms or molecules are far apart and move freely. Gases are compressible and take both the shape and volume of their container.

Classification by Composition

Matter can also be classified by its composition: elements, compounds, and mixtures.

• Pure Substance: Made up of only one component with invariant composition. Can be an element or a compound.

• Mixture: Composed of two or more components in variable proportions.

Classification of Pure Substances

• Element: Cannot be chemically broken down into simpler substances. Composed of a single type of atom (e.g., helium).

• Compound: Composed of two or more elements in fixed, definite proportions (e.g., water, sugar).

Classification of Mixtures

• Heterogeneous Mixture: Composition varies from one region to another. Different substances can be seen (e.g., salt and sand mixture).

• Homogeneous Mixture: Uniform composition throughout; appears as a single substance (e.g., sweetened tea).

Type	Definition	Example
Element	Cannot be broken down chemically	Helium
Compound	Composed of two or more elements in fixed proportions	Water (H2O)
Heterogeneous Mixture	Non-uniform composition	Salt and sand
Homogeneous Mixture	Uniform composition	Sugar dissolved in water

The Scientific Method

Empirical Approach and Steps

The scientific method is a systematic process for understanding nature through observation and experimentation.

• Observations: Data collected about the characteristics or behavior of nature.

• Hypothesis: A tentative explanation for observations; must be falsifiable.

• Experimentation: Testing hypotheses through controlled experiments.

• Law: A brief statement summarizing past observations and predicting future ones (e.g., Law of Conservation of Mass).

• Theory: A well-established explanation for observations, validated by experiments (e.g., Dalton's atomic theory).

Comparison of Law and Theory

• Law: Summarizes a series of related observations; describes what nature does.

• Theory: Explains the underlying reasons for observations; explains why nature behaves as it does.

Measurement in Chemistry

Importance of Scientific Measurement

Scientific data can be qualitative (descriptive) or quantitative (measurable). Quantitative data are objective and use standardized units, such as those in the International System of Units (SI).

• Qualitative Data: Observational, subjective (e.g., color, shape).

• Quantitative Data: Measurable, objective (e.g., mass, volume).

Early Ideas About the Building Blocks of Matter

Historical Development of Atomic Theory

The concept of atoms as the fundamental building blocks of matter dates back to ancient Greece.

• Leucippus and Democritus (5th century B.C.E.): Proposed that matter is composed of small, indestructible particles called atoms.

• Plato and Aristotle: Rejected atomic theory; believed matter was continuous and composed of fire, air, earth, and water.

• John Dalton (1766–1844): Provided experimental evidence for the existence of atoms and formulated the modern atomic theory.

Dalton's Atomic Theory and Related Laws

• Law of Conservation of Mass: Matter is neither created nor destroyed in a chemical reaction.

• Law of Definite Proportions: A compound always contains the same elements in the same proportions by mass.

• Law of Multiple Proportions: When two elements form more than one compound, the masses of one element that combine with a fixed mass of the other are in ratios of small whole numbers.

Additional info: Some content and examples were expanded for clarity and completeness based on standard General Chemistry curriculum.