Matter

Chemistry 100 Study Notes

Classification of Matter

  • Matter: Anything that has mass and takes up space.
    • Pure Substances: Have a fixed or definite composition.
    • Elements: The simplest type of pure substance containing only one type of material.
    • Compounds: Atoms of two or more elements chemically combined in the same proportion, examples include:
      • Water (H₂O)
      • Hydrogen peroxide (H₂O₂)
      • Table salt (NaCl)
      • Sugar (C₁₂H₂₂O₁₁)

Mixtures

  • Mixture: A type of matter consisting of two or more substances that are physically mixed but not chemically combined.
    • Properties:
    • Two or more substances in different proportions.
    • Substances that can be separated by physical methods.
    • Types:
    • Homogeneous Mixtures (Solutions): Uniform throughout.
    • Heterogeneous Mixtures: Do not have a uniform composition.

Organization of Matter by Composition

  • Matter is organized as follows:
    • Pure Substances
    • Elements
    • Compounds
    • Mixtures
    • Homogeneous
    • Heterogeneous
    • Examples include:
      • Copper (Element)
      • Water (Compound)
      • Brass (Mixture of copper and zinc)
      • Water and copper (Heterogeneous)

Study Check 1

  • Identify each as a pure substance or a mixture:
    • A. Pasta and tomato sauce: Mixture
    • B. Aluminum foil: Pure substance
    • C. Helium: Pure substance
    • D. Air: Mixture
  • Identify each as homogeneous or heterogeneous mixture:
    • A. Hot fudge sundae: Heterogeneous
    • B. Shampoo: Homogeneous
    • C. Sugar water: Homogeneous
    • D. Peach pie: Heterogeneous

States of Matter

  • States of Matter: Matter exists in three physical states: solid, liquid, and gas.
    • Macro Characteristics:
    • Shape
    • Volume
    • Micro Characteristics (Molecular Level):
    • Arrangement of particles
    • Interactions between particles
    • Movement of particles
Solids
  • Definite shape and volume.
  • Particles are close together in a fixed arrangement.
  • Particles move very slowly with strong attractions.
Liquids
  • Indefinite shape (takes shape of container) but definite volume.
  • Particles are close together but mobile and move with moderate speed.
Gases
  • Indefinite shape and volume (takes both from the container).
  • Particles are far apart with no attractions and move very quickly.

Study Check 2 Answers

  • Identify each as: 1) solid; 2) liquid; 3) gas:
    • a) Liquid (2): Definite volume, takes shape of the container.
    • b) Gas (3): Particles are moving rapidly.
    • c) Gas (3): Fills the volume of a container.
    • d) Solid (1): Fixed arrangement of particles.
    • e) Liquid (2): Particles close together, mobile.

Physical Properties and Changes

  • Physical Properties: Characteristics observed or measured without changing the identity of a substance. Examples include:
    • Shape
    • Physical state
    • Boiling and freezing points
    • Density
    • Color
  • Physical Changes: Changes in state or physical shape; no change in identity or composition occurs, and no new substances are produced.

Chemical Properties and Changes

  • Chemical Properties: Relate to how substances interact with others to change into new substances.
  • Chemical Changes: Original substances are transformed into one or more new substances, introducing new chemical and physical properties. Examples:
    • Silver tarnishing (reaction with air forming a black coating).
    • Wood burning (producing ash, carbon dioxide, water vapor, and heat).
    • Iron rusting (reaction with oxygen forming rust).

Study Check 3

  • Classify each:
    • Ice melts in the Sun: Physical
    • Copper is a shiny metal: Physical
    • Paper can burn: Chemical
    • A silver knife can tarnish: Chemical
    • A magnet removes iron particles from a mixture: Physical
  • Classify as Physical or Chemical Changes:
    • Burning a candle: Chemical
    • Ice melting on the street: Physical
    • Toasting a marshmallow: Chemical
    • Cutting a pizza: Physical
    • Iron rusting in an old car: Chemical

Temperature Conversions

  • Temperature: Measure of how hot or cold a substance is.
  • Two common temperature scales:
    • Celsius (°C):
    • Freezing point: 0°C
    • Boiling point: 100°C
    • Fahrenheit (°F):
    • Freezing point: 32°F
    • Boiling point: 212°F
  • Conversion equations:
    • From Celsius to Fahrenheit: T<em>F=1.8(T</em>C)+32T<em>F = 1.8(T</em>C) + 32
    • From Fahrenheit to Celsius: T<em>C=(T</em>F−32)/1.8T<em>C = (T</em>F - 32)/1.8

Temperature Conversion Calculations

  • Example Convert 40°C to Fahrenheit:
    • TF=1.8(40)+32=72+32=104°FT_F = 1.8(40) + 32 = 72 + 32 = 104°F
  • Example Convert 50°F to Celsius:
    • TC=(50−32)/1.8=(18)/1.8=10°CT_C = (50 - 32)/1.8 = (18)/1.8 = 10°C

Study Check 4

  • Given temperature conversions:
    • A) -15 °C to °F
    • B) 455 °F to Celsius

Study Check 4 Answers

  • A) State Given: -15 °C | Need: T_F | Calculation:
    • TF=1.8(−15)+32=−27+32=5°FT_F = 1.8(-15) + 32 = -27 + 32 = 5°F
  • B) State Given: 455 °F | Need: T_C | Calculation:
    • TC=(455−32)/1.8=423/1.8=235°CT_C = (455 - 32) / 1.8 = 423 / 1.8 = 235°C

Kelvin Temperature Scale

  • Kelvin: Absolute Temperature Scale.
    • Absolute Zero (0 K) = -273 °C.
    • Relationship with Celsius: T<em>K=T</em>C+273T<em>K = T</em>C + 273

Comparison of Temperature Scales

  • Reference points for boiling and freezing points of water:
    • Boiling Point:
    • 373 K = 100°C = 212°F
    • Freezing Point:
    • 273 K = 0°C = 32°F
    • Normal Body Temperature:
    • 310 K = 37°C = 98.6°F

Energy in Chemistry

  • Energy: The ability to do work. Example: Work done while climbing.
Forms of Energy
  • Kinetic Energy: Energy of motion, examples:
    • Walking up stairs
    • Water trickling down a stream
    • A fast-moving skier
    • Burning gasoline.
  • Potential Energy: Stored energy or energy of position, examples:
    • Chemical energy in batteries
    • Gasoline in a car
    • A skier at a mountain's peak.

Study Check 5

  • Identify energy as potential or kinetic:
    • Swimming: Kinetic
    • Peanut butter and jelly sandwich: Potential
    • Mowing the lawn: Kinetic
    • Gasoline in gas tank: Potential

Heat as Kinetic Energy

  • Heat: Energy associated with particles' motion.
    • Measured in units: joules (J) or calories (cal):
    • 1 calorie: Amount of energy needed to raise the temperature of 1 g of water by 1 °C.
      • 4.184J=1cal4.184 J = 1 cal
      • 1kJ=1000J1 kJ = 1000 J
      • 1kilocalorie(kcal)=1000cal1 kilocalorie (kcal) = 1000 cal

Energy and Nutrition

  • Diet provides kcal of energy.
    • Carbohydrates: Primary fuel.
    • Fats, then proteins: Secondary energy sources.
    • Energy balance affects weight (gain or loss).

Energy Values for Food

  • Energy on food labels shown as Cal (nutritional Calorie).
    • Different units globally (kJ).
    • Calorimeter: Measures the energy value by burning food samples.
    • Energy values:
    • 1Cal=1kcal=1000cal1 Cal = 1 kcal = 1000 cal
    • 1Cal=4.184kJ=4184J1 Cal = 4.184 kJ = 4184 J

Energy Value Calculation Example

  • Milk: 13 g carb, 9 g fat, 9 g protein.
  • Energy calculation for 1 cup of milk:
    • 13 g carbohydrates: 13extgimes4extkcal/g=52extkcal13 ext{ g} imes 4 ext{ kcal/g} = 52 ext{ kcal}
    • 9 g fat: 9extgimes9extkcal/g=81extkcal9 ext{ g} imes 9 ext{ kcal/g} = 81 ext{ kcal}
    • 9 g protein: 9extgimes4extkcal/g=36extkcal9 ext{ g} imes 4 ext{ kcal/g} = 36 ext{ kcal}
    • Total = 52 + 81 + 36 = 169 kcal (rounded to 170 kcal)

Study Check 6

  • Egg Composition: 6 g protein, 6 g fat, 0 g carbohydrates.
  • Energy Calculation:
    • extProtein:6extgimes4extkcal/g=24extkcalext{Protein: } 6 ext{ g} imes 4 ext{ kcal/g} = 24 ext{ kcal}
    • extFat:6extgimes9extkcal/g=54extkcalext{Fat: } 6 ext{ g} imes 9 ext{ kcal/g} = 54 ext{ kcal}
    • Total kcal = 24 + 54 + 0 = 78 kcal (80 Food Calories)

Specific Heat

  • Specific Heat (SH): Measure of heat absorption capability.
    • Defined as the amount of heat needed to raise the temperature of 1 g of a substance by 1 °C.
    • Units: J/g °C or cal/g °C.
  • Formula:
    • SH = rac{ ext{heat (J or cal)}}{ ext{grams} imes ext{ΔT}}
  • Characteristics:
    • Low SH: Efficient heat transfer (e.g., aluminum, copper).
    • High SH: Efficient heat absorption (e.g., water).

Calculating Using Specific Heat

  • Heat lost or gained determined by:
    • Mass of substance (g)
    • Temperature change (ΔT)
    • Specific Heat (SH) (J/g °C or cal/g °C).
  • Heat equation:
    • extHeat=extmass(g)imesextΔTimesextSHext{Heat} = ext{mass (g)} imes ext{ΔT} imes ext{SH}

Study Check 7

  • Copper pan: 135 g, raise temp from 26 °C to 328 °C, specific heat = 0.385 J/g °C.

Study Check 7 Answer

  • Temperature change:
    • ΔT=328°C−26°C=302°CΔT = 328°C - 26°C = 302°C
  • Heat calculation:
    • extHeat(J)=135extgimes302°Cimes0.385extJ/g°Cext{Heat (J)} = 135 ext{ g} imes 302 °C imes 0.385 ext{ J/g°C}
    • extHeat(J)=135imes302imes0.385=15.7extkJext{Heat (J)} = 135 imes 302 imes 0.385 = 15.7 ext{ kJ}

Changes of State

  • Heat is always involved in state changes:
    • Solid melts to liquid.
    • Liquid boils to gas.
    • Gas condenses to liquid.
    • Solid undergoes sublimation to gas.

Melting and Freezing

  • Melting: solid to liquid (requires heat).
  • Freezing: liquid to solid (gives off heat).
  • Temperature specific, for water:
    • Freezing/melting point: 0°C.
  • Time-dependent.

Heat of Fusion

  • Heat of Fusion: Amount of heat added for melting or removed for freezing.
    • For water: 80extcal/g80 ext{ cal/g} at 0 °C.
    • Formula for calculating Heat of Fusion:
    • extHeat=extmassimesextheatoffusionext{Heat} = ext{mass} imes ext{heat of fusion}
  • Conversion factors:
    • 80extcal/1extgH2O80 ext{ cal} / 1 ext{ g H₂O}
    • 334extJ/1extgH2O334 ext{ J} / 1 ext{ g H₂O}

Calculating Heat to Freeze Water

  • Example: 25.0 g of water at 0 °C freezes.
    • Heat lost formula:
    • extHeat=extmassimesextheatoffusionext{Heat} = ext{mass} imes ext{heat of fusion}
    • 25.0 ext{ g H₂O} imes 334 ext{ J} imes rac{1 ext{ kJ}}{1000 ext{ J}} = 8.35 ext{ kJ}

Vaporization

  • Vaporization: Liquid to gas conversion.
    • Can occur as:
    • Boiling: Gas forms throughout the liquid at its boiling point.
    • Evaporation: Can occur at any temperature at the liquid's surface.
  • Condensation: Reverse process of vaporization.

Heat of Vaporization

  • Energy to convert 1 g of liquid to gas (boiling point): 540extcal/g540 ext{ cal/g} or 2260extJ2260 ext{ J} for water.

Heating and Cooling Curves

  • Heating Curves: Diagram where diagonal lines indicate temperature changes and horizontal lines indicate state changes.

Using Heating Curves

  • For heating without change of state:
    • Use heat equation for heating: extHeat=extgimesextΔTimesextSHext{Heat} = ext{g} imes ext{ΔT} imes ext{SH}
  • For change of state:
    • Use heat of fusion or heat of vaporization:
    • extHeat=gimesHFusionext{Heat} = g imes H_{Fusion}
    • extHeat=gimesHVaporizationext{Heat} = g imes H_{Vaporization}