General Chemistry I: Introduction, Measurement, and Problem Solving

General Chemistry I: Course Overview

Introduction to CHEM 1770

This course, General Chemistry I (CHEM 1770), introduces students to the foundational principles of chemistry, including the composition, structure, properties, and transformations of matter. Chemistry is often called the "central science" because it connects and underpins fields such as medicine, agriculture, food science, and engineering.

• Instructor: Dr. Joe Race

• Lecture Times: Tuesdays and Thursdays, 2:10 pm - 3:35 pm, Troxel Hall

• Course Materials: Notes and lecture recordings are available on Canvas. Required lab course: CHEM 1770L.

• Communication: Contact via email or office hours for questions and support.

The Importance of Chemistry

Why Study Chemistry?

Chemistry is essential for understanding the world around us and is fundamental to many scientific and practical fields.

• Medicine: Understanding drug composition and biochemical processes.

• Agriculture and Food Science: Soil chemistry, fertilizers, and food preservation.

• Everyday Life: Cooking, making coffee, fueling cars, and more.

• Central Science: Chemistry bridges biology, physics, and environmental science.

Chapter E: Essentials – Units, Measurements, and Problem Solving

Learning Objectives

• Understand the scientific method and its application in chemistry.

• Use and convert between Fahrenheit, Celsius, and Kelvin temperature scales.

• Report measurements with the correct number of significant figures.

• Apply dimensional analysis to solve problems involving units, mass, and volume.

The Scientific Method

The scientific method is a systematic approach to understanding natural phenomena through observation, hypothesis formation, experimentation, and theory development.

• Observation: Gathering data about the natural world.

• Hypothesis: A tentative explanation or prediction that can be tested.

• Experiment: Testing the hypothesis under controlled conditions.

• Theory: A well-substantiated explanation of some aspect of the natural world.

• Law: A concise statement or mathematical equation describing a fundamental relationship (e.g., $P = kT$ for gas pressure and temperature).

Units and Measurements

Accurate measurement is fundamental in chemistry. The International System of Units (SI) is used for consistency.

• Base SI Units: Meter (m), kilogram (kg), second (s), mole (mol), kelvin (K), ampere (A), candela (cd).

• Metric System: Uses prefixes (e.g., milli-, centi-, kilo-) to indicate multiples or fractions of units.

Temperature Scales

• Celsius (°C): Water freezes at 0°C and boils at 100°C.

• Fahrenheit (°F): Water freezes at 32°F and boils at 212°F.

• Kelvin (K): Absolute temperature scale; 0 K is absolute zero.

Conversion Formulas:

• $F = (1.8 \times C) + 32$

• $C = (F - 32)/1.8$

• $K = C + 273.15$

• $C = K - 273.15$

Significant Figures

Significant figures (sig figs) reflect the precision of a measurement. Reporting the correct number of significant figures is essential in scientific communication.

• All nonzero digits are significant.

• Zeros between nonzero digits are significant.

• Leading zeros are not significant.

• Trailing zeros are significant only if there is a decimal point.

• Exact numbers (from counting or definitions) have infinite significant figures.

Examples:

• 0.010 mL has 2 significant figures.

• 3.40 × 102 has 3 significant figures.

• 1.00 × 103 has 3 significant figures.

• 1 inch = 2.54 cm (exact, infinite sig figs).

Rules for Calculations

• Multiplication/Division: The result has as many significant figures as the factor with the fewest significant figures.

• Addition/Subtraction: The result has as many decimal places as the quantity with the fewest decimal places.

Dimensional Analysis (Factor-Label Method)

Dimensional analysis is a systematic approach to problem solving that uses conversion factors to move from one unit to another.

• Write all units throughout the calculation.

• Set up conversion factors so that units cancel appropriately.

• Check that the final unit matches the desired quantity.

Example: Convert 60.0 mph to m/s.

• 1 mile = 1609.3 meters (inexact)

• 1 hour = 3600 seconds (exact)

Density

Density is the ratio of an object's mass to its volume. It is a physical property that can be used to identify substances.

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

• Units: g/cm3 (solids), g/mL (liquids), g/L (gases)

Example: What is the mass of 3.46 L of water? (Density = 0.997 g/mL)

• Convert L to mL: $3.46 \text{ L} \times 1000 \frac{\text{mL}}{1 \text{L}} = 3460 \text{ mL}$

• Calculate mass: $3460 \text{ mL} \times 0.997 \frac{\text{g}}{\text{mL}} = 3449 \text{ g}$

Energy and Its Forms

Types of Energy

• Kinetic Energy: Energy associated with motion.

• Potential Energy: Energy due to position or composition.

• Chemical Energy: A form of potential energy stored in chemical bonds.

• Thermal Energy: Energy associated with the temperature of an object.

Work and Heat

• Work: The result of a force acting through a distance.

• Heat: The transfer of thermal energy between objects due to temperature difference.

Energy Units and Conversion Factors

• Joule (J): SI unit of energy.

• Calorie (cal): 1 cal = 4.184 J

• Calorie (Cal, food calorie): 1 Cal = 1 kcal = 1000 cal

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

Thermodynamics: Systems and Surroundings

Definitions

• System: The part of the universe being studied (e.g., a chemical reaction).

• Surroundings: Everything outside the system.

Energy Changes

• Exothermic Process: Releases heat to the surroundings; energy of the system decreases. (Negative sign for heat change)

• Endothermic Process: Absorbs heat from the surroundings; energy of the system increases. (Positive sign for heat change)

Example: A hand warmer is an exothermic process; the system (hand warmer) releases heat to the surroundings (your hands).

Table: Comparison of Temperature Scales

Scale	Freezing Point of Water	Boiling Point of Water	Absolute Zero
Celsius (°C)	0	100	-273.15
Fahrenheit (°F)	32	212	-459.67
Kelvin (K)	273.15	373.15	0

Table: Significant Figures Rules

Rule	Example	Sig Figs
Nonzero digits	123	3
Zeros between nonzero digits	1002	4
Leading zeros	0.0025	2
Trailing zeros with decimal	2.300	4
Trailing zeros without decimal	2300	2 (ambiguous)
Exact numbers	1 inch = 2.54 cm	Infinite

Problem Solving Strategies

• Identify the information given and what is being asked.

• Develop a strategy (plan) for solving the problem.

• Carry out the calculations, keeping track of units and significant figures.

• Check the answer for reasonableness and correct units.

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