Introduction to General Chemistry Unit 1

Introduction to General Chemistry Unit 1

Matter, Measurements, and Calculations

Objective 1: Relate Chemistry to Everyday Life

• Chemistry as Central Science

  • Chemistry connects to many other scientific disciplines:

    • Nutrition

    • Health Sciences

    • Microbiology

    • Physiology

  • Chemicals are ubiquitous in all forms of matter:

    • Found in food, clothing, medicine, etc.

Chemical Applications

• Medicine

• Cosmetics

• Fuels

• Pollution Control

• Food and Additives

Objective 2: Describe the Scientific Method as Applied to Chemistry

• Steps of the Scientific Method:

  1. Observation

  2. Hypothesis formulation

  3. Experimentation to test the hypothesis

  4. Reformulate hypothesis based on experimental results

  5. Develop theories through repeated experimentation

  • Theories may evolve as new experimental data becomes available

Compare and Contrast

• Hypothesis

  • Defined as scientific guesses that are testable through experimentation

• Models

  • Representations that help explain scientific phenomena

• Theories

  • Detailed explanations based on experimental evidence; can be revised with new data

Compare and Contrast Basic and Applied Research

• Basic Research

  • Focus on acquiring knowledge for its own sake

  • Example: Investigating the properties and structure of compounds such as purines

• Applied Research

  • Focus on solving specific problems in industry or environmental contexts

  • Example: Developing less expensive methods to manufacture plastic

Objective 3a: Definitions

• Matter

  • Defined as the substance or material that comprises all physical objects; anything that possesses mass

• Chemistry

  • Defined as the scientific study of matter and its transformations

Objective 3b: Classify and Differentiate Substances

• Pure Substances

  • Composed of only one type of element or compound

• Impure Substances

  • Mixtures consisting of two or more elements or compounds

  • Heterogeneous Mixtures

    • Comprised of visibly different substances

Pure Substances

• Elements

  • Contain only one type of atom

  • Examples:

    • Monatomic: He, Ne, Na, Ag

    • Diatomic: O2, Cl2, F2, Br2, N2, I2

    • Polyatomic: S8, P4, C60

Compounds

• Defined as substances made from identical molecules of chemically bonded atoms or ions in a specific proportion.

• Examples of Compounds:

  • Water (H2O)

  • Sucrose (C12H22O11)

  • Table Salt (NaCl)

Mixtures

• Heterogeneous Mixtures

  • Different components are distinguishable as separate substances

• Homogeneous Mixtures

  • Consistent composition throughout

  • Solutions: All components in the same phase, appear clear

  • Suspensions: Components in different phases, appear cloudy

Objective 4: Compare and Contrast Physical and Chemical Properties

• Physical Properties

  • Defined as properties not involving the transformation of substances into different substances.

  • Examples: Length, volume, boiling point, mass, temperature, density, weight, phase.

  • Physical Quantities: Numerical descriptions of a substance's properties (e.g., 28.0 kg)

• Chemical Properties

  • Characterized by a substance's ability to transform its chemical identity through chemical changes.

  • Involves changes in how atoms bond with one another.

Compare and Contrast Physical and Chemical Changes

• Physical Change

  • Results in a change of physical properties without altering the chemical identity of a substance

  • Ex: Phase changes, changes in appearance, etc.

• Chemical Change (Chemical Reaction)

  • Occurs when atoms are rearranged to form a substance with a new chemical identity

Objective 5: Metric Units

• Define and approximate the magnitude of key metric units for various measurements:

  • Length: Meter (m)

  • Mass: Kilogram (kg)

  • Volume: Liter (L)

  • Energy: Joule (J)

  • Temperature: Degrees Celsius (°C) or Kelvin (K)

  • US Customary System Equivalents:

    • 1 yard = 0.914 m

    • 2.2 pounds = 1 kg

    • 1 quart = 0.946 L

    • 1 calorie = 4.184 J

    • 1°F = (°C × 1.8) + 32

Metric Prefixes

• Mega: Million (M)

• Kilo: Thousand (k)

• Deci: Tenth (d)

• Centi: Hundredth (c)

• Milli: Thousandth (m)

• Micro: Millionth (μ)

Objectives 6 and 7: Converting Units

• Become proficient in converting units within the metric system and between metric and USCS with the right equivalency values.

Converting Units Using Dimensional Analysis

• Method: Start by familiarizing with component conversions. For instance:

  • 1 ft = 12 in

• Two conversion factors arise:

  • rac{12 ext{ in}}{1 ext{ ft}}

  • rac{1 ext{ ft}}{12 ext{ in}}

Example Problem

• Convert 5 feet to inches:

  • 5 ext{ ft} imes rac{12 ext{ in}}{1 ext{ ft}} = 60 ext{ in}

Considerations in Choosing Conversion Factors

• Ensure that the chosen factors cancel all units except the target units.

• Setup: Initial unit should be positioned opposite from the final target unit to ensure proper cancellations.

Choosing the Correct Conversion Factor

• Example: Convert 50 m to centimeters

  1. Unit equivalency: 1 ext{ cm} = 0.01 ext{ m}

  2. Conversion factors can be written as rac{1 ext{ cm}}{0.01 ext{ m}} or rac{0.01 ext{ m}}{1 ext{ cm}}

  3. Choosing rac{1 ext{ cm}}{0.01 ext{ m}} ensures correct placement of units.

Completing the Problem

• Multiply given measurement by the chosen conversion factor:

  • ext{cm} = 50 ext{ m} imes rac{1 ext{ cm}}{0.01 ext{ m}} = 5000 ext{ cm}

Objective 7: Converting Between English and Metric

• Translate units between the two systems utilizing conversion tables effectively.

Example Problem

• Convert 31.0 in to cm:

  • Given: 1 ext{ in} = 2.54 ext{ cm}

  • Work: 31.0 ext{ in} imes rac{2.54 ext{ cm}}{1 ext{ in}} = 78.7 ext{ cm}

Problem Example

• Convert 55.0 mi/h to mi/min:

  • Given: 1 h = 60 ext{ min}

  • Work: 55.0 ext{ mi/h} imes rac{1 h}{60 ext{ min}} = 0.917 ext{ mi/min}

Conversion Factors in Series

• Convert 55 mi/h to m/s:

  • Utilized multiple equivalencies:

    • 1 ext{ km} = 0.62 ext{ mi}

    • 1 ext{ km} = 1000 ext{ m}

    • 1 ext{ h} = 60 ext{ min}

    • 1 ext{ min} = 60 ext{ s}

  • Final Setup:

    • 55 ext{ mi/h} imes rac{1 ext{ km}}{0.62 ext{ mi}} imes rac{1000 ext{ m}}{1 ext{ km}} imes rac{1 ext{ h}}{60 ext{ min}} imes rac{1 ext{ min}}{60 ext{ s}} = 24.6 ext{ m/s}

Conversion Among Temperature Units

• Conversion Formulas:

  • K = {}^{ ext{o}}C + 273.15

  • {}^{ ext{o}}C = K - 273.15

  • {}^{ ext{o}}C = rac{5}{9} ({}^{ ext{o}}F - 32)

  • {}^{ ext{o}}F = 1.8 imes {}^{ ext{o}}C + 32

Example Temperature Conversions

• Convert 350°F to °C and K:

  • {}^{ ext{o}}C = rac{(350-32) imes 5}{9} = 177^{ ext{o}}C

  • K = 177 + 273 = 450 K

• Convert -40°C to °F:

  • {}^{ ext{o}}F = rac{9}{5}(-40) + 32 = -40°F

• Convert 298 K to °C:

  • {}^{ ext{o}}C = 298 - 273 = 25^{ ext{o}}C

Your Turn

• Convert 578 cm to m:

• Convert 2310 g to kg:

More In-Class Problems

• Convert 22.0 in to cm

• Convert 4.78 kg to lb

Temperature Conversions

• Convert 50.0°F to °C

• Convert -10°C to °F

Kelvin Conversions

• Convert 25°C to K

• Convert 100K to °C

Objective 8: Compare and Contrast Weight, Mass, and Volume

• Volume: Amount of space occupied, measured in cubic meters (m³) or Liters (L).

  • 1 L = 1 dm³ = 1000 cm³

• Mass: Amount of material or inertia, measured in kilograms (kg), grams (g), or milligrams (mg), not dependent on gravity.

• Weight: The force of gravity on a mass, measured in pounds (lbs) or Newtons (N), varying with gravitational force.

Objective 9: Define Energy

• Defined as the capacity to do work; ability to exert a force and move an object

  • Potential Energy: Energy stored in an object

  • Kinetic Energy: Energy in a moving object

Objective 10: Measuring Devices

• Length: Meter Stick

• Mass: Balance

• Volume: Graduated Cylinder, Buret, Pipet

• Temperature: Thermometer

Types of Balances

• Top Loading Balance

• Analytical Balance

Measuring Volume

• Buret

• Gas Buret

• Volumetric Pipet

• Graduated Pipet

• Pipet Bulb

• Syringe

• Volumetric Flask

Objective 11: Utilize Scientific Notation

• Scientific notation is structured as follows:

  • X.YZ imes 10^{a}

  • Where X.YZ is the coefficient (value between 1 and 10)

  • 10^{a} indicates ten multiplied by itself a times

Examples of Powers of Ten

• For positive exponents:

  • 10^{4} = 10 imes 10 imes 10 imes 10 = 10000

• For negative exponents:

  • 10^{-4} = rac{1}{(10 imes 10 imes 10 imes 10)} = 0.0001

Objective 11 (continued)

• Example in use: Calculating a mass in grams:

  • 2.32 imes 10^{4} = 2.32 imes 10000 = 23000

  • 4.53 imes 10^{-4} = 4.53 imes 0.0001 = 0.000453

To Convert Between Notations

• From normal notation to scientific notation:

  • Move the decimal to position it between 1 and 10

  • Exponent sign corresponds to direction moved: left is positive, right is negative

• From scientific to normal notation:

  • Move decimal as per given exponent to transform back

Objective 12: Define Density and Perform Density Calculations

• Density defined by the relation:

  • D = rac{m}{v}

• Given two of the values (Density, Mass, Volume), calculate the third

Example Problems

• Density Calculation:

  • Given unknown liquid with volume 9.02 mL and mass 8.31g:

  • d = rac{m}{v} = rac{8.31g}{9.02 mL} = 0.92 ext{g/mL}

• Calculate mass of a 30 mL sample with density of 0.92 g/mL:

  • m = d imes v = (0.92 ext{ g/mL}) imes (30 ext{ mL}) = 27.6 ext{ g}

• Calculate volume of a 100g sample with density of 0.92 g/mL:

  • v = rac{m}{d} = rac{100g}{0.92 g/mL} = 108.7 mL

In-Class Density Problems

• Q1: Given empty graduated cylinder mass 22.32 g, filled mass 26.82 g, liquid volume 5.2 mL, compute density:

  • Density calculation required

• Q2: Find volume of 5.23 kg of lead:

  • Look up density of lead

• Q3: Mass of 1.5 Liters of ethyl alcohol:

  • Refer to density table for calculations