Chapter 4 Notes: Properties of Matter

Chapter 4: Properties of Matter

Overview

  • This chapter covers chemical reactions, chemical changes, physical changes, temperature change calculations, and energy.

Properties of Substances

  • Every substance possesses characteristic properties that define its unique identity.
  • These properties differentiate substances and are classified as physical or chemical.

Physical Properties

  • Inherent characteristics of a substance.
  • Do not require a chemical change to be revealed.
  • Examples:
    • Color
    • Taste (if safe)
    • Smell
    • State of matter (solid, liquid, gas)
    • Density
    • Melting point
    • Boiling point

Chemical Properties

  • Describe chemical reactivity.
  • How a substance reacts with others or decomposes.
  • Example: Chlorine gas
    • Physical properties: heavier than air, gas at room temperature, bad odor, greenish-yellow color, irritating, and dangerous.
    • Chemical property: reacts with sodium metal to form table salt (2Na(s) + Cl_2(g) \rightarrow 2NaCl(s)).

Example: Sodium and Chlorine

  • Sodium:
    • Shiny, soft metal (can be cut with a butter knife).
    • Reacts violently with water, producing hydrogen gas and flames.
  • Chlorine:
    • Heavier than air, greenish-yellow, dangerous gas.
  • Reaction:
    • Sodium reacts with chlorine to form table salt (sodium chloride).
    • Table salt is white, crystalline, and used to season food.

Physical vs. Chemical Changes

  • No two substances have identical physical and chemical properties.

Physical Changes

  • Changes in physical properties or states of matter.
  • No change in chemical composition or formula.
  • No new substances are formed.
  • Example: Sawing wood (changes shape, but remains wood).

Chemical Changes

  • New substances are formed.
  • Different chemical properties and composition (chemical formula) from the original material.
  • Example: Heating copper metal in air
    • Shiny pinkish copper turns black, forming copper(II) oxide (2Cu(s) + O_2(g) \xrightarrow{\Delta} 2CuO(s)).
    • The triangle ($\Delta$) represents heat.
  • Addition of heat energy is a clue to a chemical change but not always (e.g., melting).

Distinguishing Chemical and Physical Changes

  • Keywords for Chemical Changes: rusting, burning, reacting.
  • Keywords for Physical Changes: boiling, melting, cutting, pulverizing.

Examples

  • Combustion of gasoline: chemical change.
  • Digestion of food: chemical change (extracting nutrients and converting to waste).
  • Sawing wood: physical change.
  • Burning wood: chemical change.
  • Heating glass: physical change (changes shape, but not chemical composition).

Practice Examples

  • Grinding a rock into powder: physical change.
  • Hydrogen and oxygen react to form water: chemical change (2H2 + O2 \rightarrow 2H_2O).
  • Shovel rusting: chemical change (oxidation).
  • Acid and base reacting to form water: chemical change.
  • Burning sugar: chemical change (combustion).

Chemical Equations

  • Describe chemical changes.
  • Reactants (starting materials) on the left side.
  • Products (resulting substances) on the right side.
  • Arrow indicates a reaction produces or has occurred.
  • Physical changes often accompany chemical changes.
  • Example: Water decomposes into hydrogen and oxygen (2H2O(l) \rightarrow 2H2(g) + O_2(g)).
  • Composition (atomic/molecular) changes in a chemical reaction.
  • Equations show the change in chemical formulas.

Problem Solving in Chemistry

  • Based on the scientific method (observations, hypotheses, testing).
  • Involves measurements and units.
  • Steps:
    1. Read the Problem
    2. Identify What You're Looking For
    3. Make a Plan
    4. Do the Math
    5. Check Your Work

Energy

  • Capacity to do work.
  • Types of energy:
    • Potential energy: Stored energy or energy of position.
      • Example: A ball 20 feet above the ground has more potential energy than at 10 feet.
      • A diver on a diving board.
    • Kinetic energy: Energy of motion.
      • Example: Water released from a dam converts potential energy to kinetic energy.
  • Energy can be converted from one form to another.
  • In chemical reactions, energy is often released as heat.
  • SI unit for energy: joule (J).
  • 1 calorie (cal) is the energy required to change the temperature of 1 gram of water by 1 degree Celsius.
    • 1 cal = 4.184 J
  • Food calories are kilocalories (kcal or Cal, with a capital C).
    • 1 kcal = 1000 cal

Heat vs. Temperature

  • Different concepts.
  • Temperature: A state we measure with a thermometer.
  • Heat: Energy added to affect temperature change.
  • Example:
    • Beaker 1: 10 mL of water.
    • Beaker 2: 20 mL of water.
    • To raise both by 30°C, Beaker 2 requires twice the heat because it has more water.

Energy Problem Example

  • Combustion of methane (natural gas): CH4(g) + 2O2(g) \rightarrow CO2(g) + 2H2O(g)
  • Given: 802.5 kJ produced; convert to calories.
  • Strategy: kJ → J → calories.
  • Conversion factors:
    • 1 kJ = 1000 J
    • 1 cal = 4.184 J
  • Calculation: 802.5 kJ \times \frac{1000 J}{1 kJ} \times \frac{1 cal}{4.184 J} = 191,802.103 cal
  • With 4 significant figures: 1.918 \times 10^5 cal

Heat Capacity and Specific Heat

  • Physical property of each substance.
  • Specific heat: Amount of heat required to change the temperature of 1 gram of a substance by 1 degree Celsius.
  • Water has a much higher specific heat than most substances.
  • Key for temperature regulation in bodies of water and living things.

Metals as Conductors of Heat

  • Metals are good conductors of heat and electricity.
  • This is why metals are often used in cookware, to easily transmit heat from a stove or oven to food.

Calculation of Heat

  • Heat (Q) needed to change temperature: Q = m \times c \times \Delta T
    • m = mass
    • c = specific heat capacity
    • \Delta T = change in temperature

Problem Example 1

  • How much heat is needed to raise the temperature of 200 grams of water by 10°C?
  • Q = (200 g) \times (4.184 J/g°C) \times (10.0 °C)
  • Q = 8368 J
  • Q = 8.37 \times 10^3 J (with 3 significant figures)

Problem Example 2

  • Calculate specific heat of an unknown substance if 1638 J raises the temperature of 125 g from 25°C to 52.6°C.
  • \Delta T = 52.6 °C - 25.0 °C = 27.6 °C
  • Rearrange equation: c = \frac{Q}{m \times \Delta T}
    • Q = 1638 J
    • m = 125 g
    • \Delta T = 27.6 °C
  • c = \frac{1638 J}{(125 g) \times (27.6 °C)}
  • c = 0.475 J/g°C
  • The substance could be a metal due to its low specific heat.

Energy in Chemical Reactions

  • Chemical reactions either absorb or release energy.
  • Redox reactions can produce electrical energy (batteries).
  • Combustion releases heat and light.
  • Some chemical changes consume energy (e.g., electrolysis of water).
  • Plants use solar energy to convert CO2 into sugars (photosynthesis).

Laws of Energy

  • Law of conservation of energy: Energy cannot be created nor destroyed.
  • Energy can be lost through inefficient processes (dissipated as heat).
  • Decomposition of water absorbs energy; H2 and O2 have higher potential energy.
  • Burning hydrogen releases energy; water has lower potential energy.
  • Potential energy in molecules is held in chemical bonds; breaking bonds releases energy, forming bonds stores energy.

Energy Sources

  • Petroleum products, coal, and water sources.
  • All energy originally came from the sun.
  • Petroleum is a fossil fuel from plant materials broken down over time.
  • Hydrocarbons: Compounds containing carbon and hydrogen (e.g., gasoline, natural gas).

Common Hydrocarbons

  • Methane (CH_4)
  • Propane (C3H8) - used in fireplaces, water heaters, and stoves
  • Butane (C4H{10}) - used in lighters
  • Pentane (C5H{12}), Hexane (C6H{14}), Heptane (C7H{16}), Octane (C8H{18}) - found in gasoline and used to rate burning qualities in gasoline combustion engines
  • Kerosene - with more numbers of carbons and hydrogens

Natural Gas and Coal

  • Natural gas: mainly methane with small amounts of higher carbon fuels mixed in.
  • Coal: formed from plants stored under high pressure for many years.
  • Higher carbon content = more energy.
  • Reaction of carbon in coal with oxygen produces CO2 and energy.

Renewable Energy Sources

  • Solar, nuclear, biomass (ethanol), wind, synthetic fuels.
  • All sources have costs and drawbacks (economic, environmental, and efficiency considerations).
  • Need to understand numbers around input and use of different energy sources.

Chapter Summary

  • Properties of substances (physical and chemical changes).
  • Problem-solving with energy.
  • Definitions of kinetic and potential energy.
  • Heat and quantitative measurement of heat.
  • SI unit for energy.
  • Solving specific heat problems (Q = mcΔT).
  • Energy flow in chemical changes and law of conservation of energy.
  • Links to real-world applications.