Introduction to General Chemistry: Matter, Measurement, and Problem Solving

Introduction to General Chemistry

Course Overview and Organization

• Course Title: General Chemistry for Engineers (CHEN 1201)

• Instructor: Carolyn Kohlmeier, Department of Chemical and Biological Engineering

• Course Structure: Combination of lectures, recitations, homework, quizzes, and exams. Emphasis on both in-person and remote learning via Canvas.

• Grading Scheme:

  • Recitation: 10%

  • Homework: 15%

  • Online reading quizzes: 5%

  • Clickers: 5%

  • Midterm Exams (3): 15% each

  • Final Exam: 20%

• Course Communication: All logistical and administrative questions should be directed to the course email or Canvas site.

Matter, Measurement, and Problem Solving

Physical and Chemical Changes

Understanding the distinction between physical and chemical changes is fundamental in chemistry.

• Physical Change: Alters only the state or appearance of a substance, not its chemical composition.

• Chemical Change: Alters the composition of matter, resulting in the formation of one or more new substances.

• Example: Freezing water is a physical change; metabolism of food is a chemical change.

Properties of Matter

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

• Intensive Properties: Independent of the amount of substance (e.g., density, melting point).

• Example: Volume is extensive; density is intensive.

Classification of Matter

Matter can be classified based on the composition and uniformity of its constituent particles.

• Pure Substances: Constituent particles are identical.

  • Elements: Contain only one type of atom (e.g., O2).

  • Compounds: Contain molecules or ions with multiple types of atoms (e.g., H2O).

• Mixtures: Contain multiple types of particles.

  • Heterogeneous Mixtures: Particles are not uniformly mixed; can be physically separated (e.g., sand in water).

  • Homogeneous Mixtures (Solutions): Particles are uniformly mixed; transparent and chemically separable (e.g., saltwater).

SI Units and Derived Units

The International System of Units (SI) is used for scientific measurements. It includes both fundamental and derived units.

Physical Quantity	Name of Unit	Abbreviation
Mass	kilogram	kg
Length	meter	m
Temperature	kelvin	K
Amount of substance	mole	mol
Time	second	s
Electric current	ampere	A
Luminous intensity	candela	cd

Quantity	Definition	Derived Unit (Name)
Area	Length × Length	m2
Volume	Area × Length	m3
Density	Mass per unit volume	kg/m3
Speed	Distance per unit time	m/s
Acceleration	Change in speed per unit time	m/s2
Force	Mass × acceleration	kg·m/s2 (Newton, N)
Pressure	Force per unit area	kg/(m·s2) (Pascal, Pa)
Energy	Force × distance	kg·m2/s2 (Joule, J)

Metric Prefixes

Factor	Prefix	Symbol
1,000,000,000 = 109	giga	G
1,000,000 = 106	mega	M
1,000 = 103	kilo	k
100 = 102	hecto	h
10 = 101	deka	da
0.1 = 10-1	deci	d
0.01 = 10-2	centi	c
0.001 = 10-3	milli	m
0.000001 = 10-6	micro	μ
0.000000001 = 10-9	nano	n
0.000000000001 = 10-12	pico	p

Length Scales of Common Objects

• Atoms: $10^{-10}$ m

• Molecules/nanotech: $10^{-9}$ m

• Visible light: $10^{-6}$ m

• Cells: $10^{-5}$ m

• Lab glassware: $10^{-2}$ m

• Industrial process: $10^{2}$ m

Dimensional Analysis

Dimensional analysis is a method for converting between units using conversion factors.

• Example Problem: Calculate the volume occupied by 10.0 kg of titanium (Ti) with a density of 4.50 g/cm3.

Solution:

• Given: mass = 10.0 kg, density = 4.50 g/cm3

• Convert mass to grams: $10.0\ \text{kg} \times \frac{1000\ \text{g}}{1\ \text{kg}} = 10,000\ \text{g}$

• Calculate volume: $\text{Volume} = \frac{\text{mass}}{\text{density}} = \frac{10,000\ \text{g}}{4.50\ \text{g/cm}^3} = 2222.2\ \text{cm}^3$

• Convert cm3 to L: $2222.2\ \text{cm}^3 \times \frac{1\ \text{L}}{1000\ \text{cm}^3} = 2.22\ \text{L}$

Classification of Matter: Diagram

• Matter: Macroscopic substances composed of tiny "particles".

• Mixtures: Multiple types of particles.

  • Heterogeneous: Not intimately mixed, opaque, physically separable.

  • Homogeneous (Solutions): Randomly mixed, transparent, chemically separable.

• Pure Substances: Constituent particles identical.

  • Elements: Only one type of atom; particles may be atoms or molecules.

  • Compounds: Molecules/ions with multiple types of atoms.

Visual Representation: Mixtures of Gases

• Red spheres represent oxygen atoms; white spheres represent hydrogen atoms.

• A mixture of hydrogen and oxygen gas contains both H2 and O2 molecules, not chemically bonded together.

Practice Questions

• Which of the following represents a chemical change?

  • A. Freezing water to make ice cubes (Physical change)

  • B. Metabolism of food in the body (Chemical change)

  • C. Dry ice evaporating at room temperature (Physical change)

  • D. Dissolving sugar in hot coffee (Physical change)

  • E. Crushing an aluminum can (Physical change)

• Identify the extensive property:

  • A. It is a solid at 25°C. (Intensive)

  • B. It has a density of 1.38 g/cm3. (Intensive)

  • C. It melts at 62.0°C. (Intensive)

  • D. It has a volume of 0.52 cm3. (Extensive)

  • E. It is shiny. (Intensive)

Additional info: These notes cover foundational concepts in general chemistry, including the classification of matter, physical and chemical changes, properties of matter, SI units, metric prefixes, dimensional analysis, and basic problem-solving strategies. These are essential for understanding subsequent topics in chemistry.