General Chemistry I - CHM 101: Introduction and Chapter 1 Study Notes

General Chemistry I - CHM 101

Course Overview

This course introduces the fundamental principles of chemistry, focusing on the composition, structure, properties, and changes of matter. Chemistry is often referred to as the central science because it connects physical sciences with life sciences and applied sciences such as medicine and engineering.

• Instructor: Dr. Kathryn Sarachan

• Contact: [e-mail]

• Class Time: MWF 10:00 – 10:50 AM

• Location: SC 114 (Brooks Auditorium)

• Office Hours: T 12:45-2:15 pm, Th 12:30-1:30 pm, and by appointment

Chemistry: The Central Science

What is Chemistry?

Chemistry is the study of the composition, structure, properties, and behavior of matter. It is called the central science because it bridges other natural sciences, including physics, biology, and geology.

• Energy: Chemistry explains how solar panels work by understanding the properties of silicon and its ability to convert sunlight into electricity.

• Biochemistry: Chemical reactions in living organisms, such as the light produced by fireflies, are studied in biochemistry.

• Technology: The development of LEDs (light emitting diodes) relies on the chemical properties of elements like gallium, arsenic, and phosphorus.

Chapter 1: Matter, Energy & Measurement

Learning Objectives

• Classify matter by composition and state

• Explain how chemical and physical properties allow matter to be described and classified

• Define kinetic and potential energy in chemical systems

• Describe the metric system and the use of significant figures in reporting measurements

• Convert measurements correctly using dimensional analysis

Matter: Classification by Composition and State

Matter is anything that has mass and occupies space. It can be classified by its composition (what it is made of) and by its physical state (solid, liquid, or gas).

• Substance: Matter with a fixed composition and distinct properties. Substances can be elements or compounds.

• Mixture: A combination of two or more substances where each retains its own properties. Mixtures can be homogeneous (uniform throughout, also called solutions) or heterogeneous (not uniform).

Types of Substances

• Element: A substance that cannot be decomposed into simpler substances. Elements are made of identical atoms.

• Compound: A substance composed of two or more elements chemically combined in fixed proportions. Compounds have molecules containing atoms of different elements.

Law of Constant Composition (Law of Definite Proportions): The elemental composition of a compound is always the same, regardless of its source.

Mixtures

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

• Heterogeneous Mixture: Non-uniform composition (e.g., potting soil).

States of Matter

Matter exists in three primary states, each with distinct physical properties:

• Gas: No fixed shape or volume; fills the container; particles are far apart and move rapidly.

• Liquid: Fixed volume but no fixed shape; takes the shape of the container; particles are close together but can move past one another.

• Solid: Definite shape and volume; particles are closely packed in fixed positions.

Properties of Matter

• Physical Properties: Can be observed without changing the substance into another substance (e.g., boiling point, density, mass, volume, odor, hardness).

• Chemical Properties: Can only be observed when a substance is changed into another substance (e.g., flammability, reactivity with acid).

Changes in Matter

• Physical Change: Changes that do not alter the composition of a substance (e.g., changes of state, temperature, volume).

• Chemical Change: Changes that result in the formation of new substances (e.g., combustion, oxidation, decomposition).

Example: Melting ice is a physical change; burning hydrogen in oxygen to form water is a chemical change.

Extensive vs. Intensive Properties

• Intensive Properties: Do not depend on the amount of substance (e.g., color, density, boiling point).

• Extensive Properties: Depend on the amount of substance (e.g., mass, volume, energy).

Energy in Chemical Processes

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

• Kinetic Energy: Energy of motion. $KE = \frac{1}{2}mv^2$ (where $m$ is mass and $v$ is velocity).

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

Measurement and the Metric System

• SI Base Units: The International System of Units (SI) uses standard units for scientific measurements.

Physical Quantity	Name of Unit	Abbreviation
Mass	Kilogram	kg
Length	Meter	m
Time	Second	s
Temperature	Kelvin	K
Amount of substance	Mole	mol
Electric current	Ampere	A
Luminous intensity	Candela	cd

Metric Prefixes

Prefix	Abbreviation	Factor
Peta	P	$10^{15}$
Tera	T	$10^{12}$
Giga	G	$10^{9}$
Mega	M	$10^{6}$
Kilo	k	$10^{3}$
Deci	d	$10^{-1}$
Centi	c	$10^{-2}$
Milli	m	$10^{-3}$
Micro	μ	$10^{-6}$
Nano	n	$10^{-9}$
Pico	p	$10^{-12}$
Femto	f	$10^{-15}$
Atto	a	$10^{-18}$
Zepto	z	$10^{-21}$

Temperature Scales

• Celsius (°C): Based on the properties of water (0°C = freezing, 100°C = boiling).

• Kelvin (K): SI unit of temperature; absolute zero (0 K) is the lowest possible temperature.

• Fahrenheit (°F): Commonly used in the United States.

Conversion formulas:

• $K = ^\circ C + 273.15$

• $^\circ F = \frac{9}{5}(^\circ C) + 32$

Volume and Energy Units

• Volume: Derived from length; common units are liter (L) and milliliter (mL). $1\,L = 1\,dm^3 = 1000\,cm^3$

• Energy: SI unit is the joule (J); $1\,J = 1\,kg\,m^2/s^2$. The calorie (cal) is also used: $1\,cal = 4.184\,J$

Uncertainty in Measurement and Significant Figures

• Exact Numbers: Counted or defined values (e.g., 12 eggs in a dozen).

• Inexact Numbers: Measured values, subject to uncertainty due to instrument limitations.

• Accuracy: How close a measurement is to the true value.

• Precision: How close repeated measurements are to each other.

Significant Figures

• All nonzero digits are significant.

• Leading zeros are not significant.

• Captive zeros (between nonzero digits) are significant.

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

• Exact numbers have an infinite number of significant figures.

Examples:

• 3456 (4 sig figs)

• 0.0486 (3 sig figs)

• 16.07 (4 sig figs)

• 9.300 (4 sig figs), 9300 (2 sig figs)

Significant Figures in Calculations

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

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

Example: $6.38 \times 2.0 = 13$ (2 sig figs)

Example: $8 + 11.934 = 19.934 \rightarrow 19.9$ (3 sig figs)

Dimensional Analysis (Unit Conversions)

Dimensional analysis is a method for converting one unit to another using conversion factors.

• Identify the conversion factor(s) needed.

• Set up the calculation so that units cancel appropriately.

Example: Convert 3.000 m to cm:

• Conversion factor: $1\,m = 100\,cm$

• Calculation: $3.000\,m \times \frac{100\,cm}{1\,m} = 300.0\,cm$

Example: Convert 3.000 m to inches:

• Use $1\,in = 2.54\,cm$ and $1\,m = 100\,cm$

• Calculation: $3.000\,m \times \frac{100\,cm}{1\,m} \times \frac{1\,in}{2.54\,cm} = 118.1\,in$

Common Conversion Factors

Quantity	Conversion
Length	1 in = 2.54 cm (exact), 1 mi = 1.6093 km, 1 m = 1.0936 yd
Mass	1 kg = 2.2046 lb, 1 lb = 16 oz
Volume	1 L = 1000 cm3, 1 gal = 4 qt, 1 qt = 0.946 L

Example: Convert 12.0 in3 to mL:

• 1 in3 = 16.3871 cm3

• 1 cm3 = 1 mL

• Calculation: $12.0\,in^3 \times \frac{16.3871\,cm^3}{1\,in^3} \times \frac{1\,mL}{1\,cm^3} = 196.6\,mL$

Additional info: These notes are based on introductory material and Chapter 1 objectives for a General Chemistry I course, using "Chemistry: The Central Science" as the primary textbook.