Notes on Introduction to Chemistry

Introduction to Chemistry

• Course Information: CHM 1025C, Spring 2025, Dr. Inga B. Pinnix
• Contact: ipinnix@fscj.edu

Course Overview

• Agenda:
  • Course Orientation
  • Syllabus discussion
  • Grading and Assignments
  • Smarter Proctoring overview
  • Module 1 details

Basic Concepts in Chemistry

• Atoms in Soda Pop:
  • Soda contains approximately 1 billion trillion atoms ($1 imes 10^{21}$ atoms in a drop).
  • Composed of various substances, primarily carbon dioxide and water.

The Nature of Chemistry

• Definition: Chemistry studies matter's composition and properties.
• Branches of Chemistry:
  • Organic Chemistry: Focus on carbon-containing compounds.
  • Inorganic Chemistry: Study of substances without carbon.
  • Biochemistry: Focus on substances derived from living organisms.
  • Green Chemistry: Design of processes to be more environmentally friendly.

Scientific Method in Chemistry

• Observations: Measure or observe aspects of nature.
• Hypothesis: Tentative interpretations of observations.
• Laws: Summary of consistent observations.
• Theories: Models explaining observations and laws with underlying causes.

Classification of Statements

• Examples:
  • (a) Metal burning in a closed container - Law
  • (b) Matter made of atoms - Theory
  • (c) Matter conserved in reactions - Law
  • (d) Wood burning in a closed container - Law

Measurement and Problem Solving

• Measurement: Precision indicated by the last reported digit. Example: Average global temperature rose by $0.6 °C ext{ (} ext{ uncertainty of } ext{ ± } 0.1 °C ext{)}$.
• Significant Figures:
  • Non-zero digits are significant.
  • Zeros between non-zero digits are significant.
  • Trailing zeros in decimal numbers count as significant.
  • Leading zeros are not significant.
• Operations:
  • Significant figures in multiplication/division = fewest significant figures.
  • Addition/Subtraction requires fewest decimal places.

SI Units and Metric System

• International System of Units (SI)
  • Base Units: meter (m), kilogram (kg), second (s), etc.
  • Metric Symbols: km, mg, mL, ns, etc.
• SI Prefixes:
  • tera (T) = $10^{12}$, giga (G) = $10^{9}$, mega (M) = $10^{6}$, kilo (k) = $10^{3}$,
  • milli (m) = $10^{-3}$, micro (μ) = $10^{-6}$, nano (n) = $10^{-9}$, pico (p) = $10^{-12}$.

Unit Conversion and Problem Solving

• Unit Factors: Ratio of two equivalents, e.g., $1 ext{ m} = 100 ext{ cm}$.
• Unit Analysis Method:
  1. Identify the unit sought.
  2. Determine the related given value.
  3. Apply unit factors for conversion.

Example Problem: Mass Conversion

• Calculate mass of a 325 mg aspirin tablet in grams:
  • Convert: $325 ext{ mg} imes \frac{1 ext{ g}}{1000 ext{ mg}} = 0.325 ext{ g}$.

The Percent Concept

• Percentage: Ratio of a part to a whole expressed out of 100.
  • Formula: % = (part/whole) × 100
• Example:
  • In bronze, if 79.2 g is copper and 10.8 g is tin, percent copper = $ rac{79.2 ext{ g}}{79.2 + 10.8 ext{ g}} imes 100 ext{= 88.0 ext{%}}$.

Volume and Density Concepts

• Volume by Displacement: Method to measure the volume of irregular objects via displaced water.
• Density: Defined as density = $rac{ ext{mass}}{ ext{volume}}$.

Temperature and Heat

• Temperature: Measure of average kinetic energy of particles.
  • Scales: Fahrenheit, Celsius, Kelvin (absolute scale).
• Heat: Total energy measure in systems, expressed in joules (J) or calories (cal).
• Specific Heat: Amount of heat required to raise 1 g of substance by $1 °C$; unit: cal/g°C.