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Page 1: Introduction

• Course: Organic and Biochemistry (CHE 124)

• Reading Assignment: General, Organic, and Biological Chemistry: An Integrated Approach, 4th Ed., Raymond.

• Chapter 1 covers Science and Measurements.

• Answers to odd-numbered problems available in the book’s index.

Page 2: What is Chemistry?

• Chemistry is defined as the study of matter and the changes it undergoes (chemical reactions).

• Emphasizes that chemistry is the central science, connecting biology, physics, engineering, medicine, and more.

Page 3: The Scientific Method

• Scientific Method: a systematic approach to gathering and interpreting information.

• Key Concepts:

  • Hypothesis: Tentative explanation or educated guess for observations.

  • Theory: Well-tested explanation for a behavior based on experimentation.

  • Law: Describes consistent and reproducible observations based on extensive testing.

Page 4: Steps of the Scientific Method

• Procedure includes:

  • Observation

  • Formulation and discarding hypotheses

  • Conducting experiments

  • Accepting or revising hypotheses to form a theory

  • Incorporating creativity in the process.

Page 5: What is Matter?

• Matter: Anything that has mass and occupies space.

• Weight: Measure of gravitational pull on matter.

• Mass: Measure of the quantity of material.

• Phases of Matter:

  • Solids: Fixed volume and shape.

  • Liquids: Fixed volume, indefinite shape, takes shape of the container.

  • Gases: Indefinite shape and volume, takes shape and volume of the container.

Page 6: Properties of Substances

• Each pure substance has unique properties:

  • Chemical Properties: Require a chemical change (e.g., flammability).

  • Physical Properties: Can be observed without changing chemical composition (e.g., boiling point, melting point, density).

• Properties classification:

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

  • Extensive Properties: Dependent on amount (e.g., mass, volume).

Page 7: What is Energy?

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

• Types of energy:

  • Potential Energy: Stored energy.

  • Kinetic Energy: Energy of motion.

Page 8: Measurements in Chemistry

• Chemistry relies on quantitative measurements.

• SI Units (International System of Units):

  • Commonly referred to as the metric system.

  • Based on decimal (powers of ten).

  • Common units include kg, L, K, °C.

• The English system is primarily used in the United States.

Page 9: SI Units Overview

• Revision of SI units approved in November 2018 based on physical constants:

  1. Length: Meter (m) - defined by speed of light.

  2. Time: Second (s) - defined by cesium-133 atom transitions.

  3. Amount of Substance: Mole (mol).

  4. Voltage: Ampere (A).

  5. Temperature: Kelvin (K) - defined by Boltzmann's constant.

  6. Luminous Intensity: Candela (cd).

  7. Mass: Kilogram (kg) - defined by Planck’s constant.

Page 10: Base SI Units

• Length (m): Distance light travels in vacuum in 1/299,792,458 seconds.

• Time (s): Defined by the fixed value of the cesium frequency.

• Mole (mol): Contains 6.022 x 10^23 elementary entities.

Page 11: More SI Units

• Ampere (A): Defined as one coulomb per second.

• Kelvin (K): Established by Boltzmann constant.

• Candela (cd): Defined by the luminous efficacy of a specific frequency.

• Kilogram (kg): Defined by Planck constant.

Page 12: Derived SI Units

• Common derived quantities and their units:

  • Area: Square meter (m²).

  • Volume: Cubic meter (m³).

  • Speed/Velocity: Meter per second (m/s).

  • Acceleration: Meter per second squared (m/s²).

Page 13: Measuring Length

• SI unit for Length: Meter (m), equivalent to 39.37 inches.

• Instruments for measuring length include meter sticks and micrometers.

Page 14: Measuring Volume

• SI Unit for Volume: Cubic meter (m³).

• Unit conversions:

  • 1000 L = 1 m³.

  • 1 cm³ = 1 mL.

• Instruments for measuring volume include graduated cylinders and pipets.

Page 15: Measuring Mass

• SI unit for mass: Kilogram (kg) based on Planck’s constant.

• Conversions: 1 kg = 1000 g; 1 g = 1000 mg.

• Instruments used: balance and scale.

Page 16: Mass versus Weight

• Mass: Amount of matter in an object.

• Weight: Gravitational force acting on that mass.

Page 17: Metric Prefixes

• Common metric prefixes:

  • Mega (M): 1,000,000 (10^6)

  • Kilo (k): 1,000 (10^3)

  • Milli (m): 0.001 (10^-3)

  • Nano (n): 0.000000001 (10^-9)

Page 18: English Conversions (Know!)

• Length Conversions:

  • 1 mile (m) = 5280 feet (ft.)

• Volume Conversions:

  • 1 gallon = 4 quarts.

• Mass Conversions:

  • 1 ton = 2000 pounds (lbs.).

Page 19: English to Metric Conversions (Know!)

• Length Conversions:

  • 1 inch = 2.54 cm.

  • 1 mile = 1.609 km.

• Volume Conversions:

  • 1 gallon = 3.785 L.

• Mass Conversions:

  • 1 lb = 453.6 g.

Page 20: Dimensional Analysis

• Dimensional analysis helps solve problems by ensuring the correct units are used.

• Conversion factors facilitate unit changes and maintain equivalent values.

Page 21: Typical Conversion Problems

• Problems may involve converting volumes (e.g., tablespoons to mL) or distances (e.g., miles to cm).

Page 22: Measuring Temperature

• Temperature scales: Celsius (°C), Fahrenheit (°F), Kelvin (K).

• Instruments: mercury thermometers, digital thermometers.

• Key temperature points:

  • Water boils at 100°C/212°F/373.15K.

  • Water freezes at 0°C/32°F/273.15K.

Page 23: Converting Temperature Scales

• Know conversion points between Fahrenheit and Celsius (32°F = 0°C) and Celsius and Kelvin.

Page 24: Typical Temperature Problems

• Problems may involve converting temperatures across different scales.

Page 25: Scientific Notation

• Used for very large/small numbers.

• Example: Avogadro's number (6.022 x 10²³).

Page 26: Accuracy vs Precision

• Accuracy: Closeness to the true value.

• Precision: Closeness of repeated measurements to one another.

Page 27: Uncertainties in Measurements

• Different measuring tools yield varying uncertainties.

Page 28: Significant Figures

• Significant figures indicate precision in measurement.

• Examples:

  • 8.00 mL has three significant figures.

  • 8 mL has one significant figure.

Page 29: Ambiguity in Significant Figures

• Using scientific notation can clarify significant figures.

Page 30: Rounding

• Rounding rules based on the first digit discarded.

Page 31: Significant Figures in Calculations

• Addition/Subtraction: Match number of decimal places.

• Multiplication/Division: Round to the number of significant figures in the least precise measurement.

Page 32: Examples of Conversions

• Problems may involve converting between different units of measure (mL to L, pounds to grams, etc.).

Page 34: Density

• Density = mass/volume (g/mL).

• Water density is 1 g/mL; temperature affects density.

Page 35: Density Problems

• Example problem involved calculating the density of a new material.

Page 36: Specific Gravity

• Compares the density of a substance with that of water, often used in various applications.

Page 38: Specific Heat

• Relates the energy absorbed or released by a mass of material as temperature changes.

Page 39: Derived SI Units with Special Names

• Commonly used derived SI units include:

  • Newton (N) for force.

  • Pascal (Pa) for pressure.

  • Joule (J) for energy/watt.

  • Degree Celsius (°C) for temperature.