What Is Matter? Dimensional Analysis

CHAPTER 1: MATTER, MEASUREMENTS, AND CALCULATIONS

1.1 WHAT IS MATTER?

  • Learning Objective:

    • Explain what matter is.

Definition of Matter

  • Matter is defined as the substance of everything around us. It includes:

    • Computers

    • Textbooks

    • Food

    • Humans

  • General Characteristics:

    • Anything that has mass and occupies space is classified as matter.

    • Mass is a key measurement that reflects the quantity of matter in an object.

Difference Between Mass and Weight

  • Common Misconception: Mass and weight are NOT the same.

    • Mass:

    • A measurement of the amount (quantity) of matter in an object.

    • Units: gram (g).

    • Example: A soccer ball and a bowling ball are similar in size, but the bowling ball has considerably more mass than the soccer ball.

    • Weight:

    • A measurement of the gravitational force acting on an object.

    • Units: pound (lb).

    • Example: A dumbbell has a weight of 5 pounds and a mass of 2.3 kg on Earth. On the Moon (where gravitational pull is 1/6 that of Earth), the dumbbell's weight is 0.8 lbs.

    • Question: What is the dumbbell's mass on the Moon?

Answer to Dumbbell's Mass Question

  • The MASS of an object remains constant regardless of location. Thus, the mass of the dumbbell is 2.3 kg both on Earth and on the Moon.

1.2 PROPERTIES AND CHANGES

  • Learning Objective:

    • Explain the difference between physical and chemical properties of matter and the changes they undergo.

Properties as Characteristics

  • In chemistry, properties are characteristics like:

    • Shape

    • Color

    • Size

    • Smell

  • Two Types of Properties:

    • Physical Properties:

    • Observable or measurable without altering the composition of matter.

    • Observation Example: You can observe the color and size of a sheet of paper without changing its composition.

    • Chemical Properties:

    • Demonstrated when attempts are made to change a substance into different kinds of matter (changing its composition).

Changes in Matter

  • Definitions:

    • Physical Changes:

    • Changes that occur without altering the composition of matter.

    • Chemical Changes:

    • Changes that involve transformations affecting composition.

Examples of Changes

  • Chemical Change Example:

    • Igniting magnesium: burns to produce white ash of magnesium oxide, altering the composition.

  • Physical Change Example:

    • Phase changes of water (ice, water, steam): while heat is added to cause each change, the composition (H2O) remains the same.

Learning Check 1.1

  • Task: Classify properties or changes as physical or chemical:

    • Milk sours (Chemical)

    • A wet handkerchief dries (Physical)

    • Fruit ripens (Chemical)

    • A stick of dynamite explodes (Chemical)

    • Air is compressed into a steel container (Physical)

    • Water boils (Physical)

Answers to Learning Check 1.1

  • Chemical Changes:

    • Milk sours: changes in taste and odor indicate new substances have formed.

    • Fruit ripens: changes in taste and odor indicate new substances formed.

    • Stick of dynamite explodes: gases and smoke indicate new substances formed.

  • Physical Changes:

    • Wet handkerchief dries: water evaporated is still water.

    • Air compressed: air is still air.

    • Water boils: phase change does not change composition.

1.3 A MODEL OF MATTER

  • Learning Objective:

    • Describe matter according to accepted scientific models.

Historical Context of Scientific Models

  • Scientific models have long been used to explain observed natural behaviors.

  • Focus on fundamental models in chemistry: Atoms and Molecules.

Molecules

  • Definition: A molecule is the smallest particle of a pure substance that retains the properties of that substance and can exist independently.

  • Example: Nitrogen exists as diatomic molecules (N₂). If divided, would eventually reach a single molecule.

Atoms

  • Definition: An atom is the smallest unit of matter that can be formed through chemical change.

  • Atoms represent the limit of chemical subdivision.

Classifying Molecules

  • Types of Molecules:

    • Diatomic Molecules: Two atoms (e.g., O₂).

    • Homoatomic Molecules: Molecules containing only one type of atom (e.g., O₂).

    • Heteroatomic Molecules: Molecules with two or more types of atoms (e.g., CO).

    • Triatomic Molecules: Three atoms (e.g., CO₂ is triatomic and heteroatomic).

    • Polyatomic Molecules: More than three atoms.

Learning Check 1.2

  • Classification Task: Classify the following:

    • Water (H₂O): Triatomic and heteroatomic.

    • Ozone (O₃): Triatomic and homoatomic.

    • Methane (CH₄): Polyatomic and heteroatomic.

Answers to Learning Check 1.2

  • Water (H₂O): Triatomic and heteroatomic.

  • Ozone (O₃): Triatomic and homoatomic.

  • Methane (CH₄): Polyatomic and heteroatomic.

1.4 CLASSIFYING MATTER

  • Learning Objective:

    • Classify matter based on observations or provided information.

Pure vs. Mixed Substances

  • Pure Substance: Matter with a constant composition and fixed set of properties. Cannot be physically separated into simpler substances.

  • Mixture: A combination of matter that can be physically separated into its components. Its composition can vary.

Types of Mixtures

  • Homogeneous Mixtures: Uniform appearance and properties throughout (e.g., sugar water).

  • Heterogeneous Mixtures: Non-uniform, varying appearances (e.g., pizza).

Classification of Substances

  • Elements: Pure substances made from one kind of atom (e.g., Oxygen O₂, Carbon C).

  • Compounds: Substances made of two or more kinds of atoms (e.g., CO, H₂O, C₁₂H₂₂O₁₁).

  • Note: Elements cannot be subdivided into simpler substances; compounds can.

Classification Summary

  • Complete Classification Scheme:

    • Matter

    • Pure Substance

      • Element

      • Compound

    • Mixture

      • Heterogeneous Mixture

      • Homogeneous Mixture

1.9 USING UNITS IN CALCULATIONS

  • Learning Objective:

    • Use the factor-unit method for solving numerical problems.

Factor-Unit Method / Dimensional Analysis

  • Concept: Set up calculations using unit factors (relationship between two values).

  • **Methodology:

    What we are solving for = What is Given × Unit Factor**

  • Approach with attention to units to cancel and arrive at the answer.

Worked Example

  • Example Problem: Convert 6 feet to inches.

    • Given: 6 feet

    • Relation: 12 inches = 1 foot

    • Set-Up:

    • h=6extfeetimesrac12extinches1extfooth = 6 ext{ feet} imes rac{12 ext{ inches}}{1 ext{ foot}}

    • Cancel units:

    • h=6imes12=72extinchesh = 6 imes 12 = 72 ext{ inches}

Practice

  • Work on Example problems and Learning Checks in the textbook!

1.10 CALCULATING PERCENTAGES

  • Learning Objective:

    • Perform calculations involving percentages.

Percentage Calculation

  • Formula: Percentage = (Part / Whole) × 100.

  • Familiarity with this calculation is vital for science applications.

1.11 DENSITY

  • Learning Objective:

    • Perform calculations involving density.

Density Definition

  • Definition: Density is a physical property equal to the mass of a sample divided by its volume.

  • Density is an intensive property, independent of the amount of substance.

  • Formula:

    • extDensity=racextMassextVolumeext{Density} = rac{ ext{Mass}}{ ext{Volume}}

Learning Check 1.19

  • Task: Calculate:

    • Mass of aluminum with 60.0 cm³ volume, density = 2.7 g/cm³.

    • extMass=60.0extcm3imesrac2.7extg1extcm3=162extgext{Mass} = 60.0 ext{ cm}^3 imes rac{2.7 ext{ g}}{1 ext{ cm}^3} = 162 ext{ g}

    • Volume of aluminum with mass of 98.5 g.

    • extVolume=98.5extgimesrac1extcm32.7extg=36.48extcm3ext{Volume} = 98.5 ext{ g} imes rac{1 ext{ cm}^3}{2.7 ext{ g}} = 36.48 ext{ cm}^3

Conclusion

  • Progress through end-of-chapter exercises for building understanding and efficiency in calculations for homework and exams!