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: $325 ext{ mg} imes rac{1 ext{ g}}{1000 ext{ mg}} = 0.325 ext{ g}$.

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.