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Transcript

  • 1. Classification of Matter
  • 2. Material is used to refer to a specific kind of matter.
    • Examples
    Wood Steel Air Milk
  • 3. A mixture is matter that contains more than one material.
    • Examples
    Granite Milk Air
  • 4. Heterogeneous Materials Mixtures in which the materials are not uniform are called heterogeneous .
  • 5. Granite
    • Granite is heterogeneous because it is composed of several different minerals. If it were crushed, you could pick out sand-sized particles of quartz, biotite, and feldspar.
  • 6. Milk
    • Appears uniform, but it can be separated into visible parts.
  • 7. Phase
    • Each separate part of a material is called a phase.
    • OR, it is any region with a uniform set of properties.
    Each material in granite is a phase. In ice water, ice & water are different phases.
  • 8. Interfaces
    • The different phases in a heterogeneous mixture are separated from each other by definite boundaries called interfaces .
    Ice Water The surfaces of the ice and water are interfaces.
  • 9. Homogeneous Materials Materials that consist of only one phase are called homogeneous .
  • 10. Examples Sugar Salt Seawater Air
  • 11. Solution
    • Only one kind of homogeneous material can be classified as a mixture, a solution .
    • Solutions are composed of more than one material:
      • Solute – dissolved material
      • Solvent – dissolving material
    • Solute particles are dispersed among the solvent uniformly.
  • 12. Example
    • Sugar Water
    Solute? Sugar Solvent? Water
  • 13. Not all solutions are liquids! Air Made of nitrogen, oxygen, and other gases…
  • 14. Molarity
    • Solutions can be different concentrations.
    • The letter “M” is used to represent the term molarity.
    • Molarity is the amount of solute in a given amount of solvent.
      • A 6M (6 molar) solution contains 6 times as much solute as 1M (1 molar) solution of the same volume.
    • Concentrated solutions have a higher ratio of solute to solvent than dilute solutions.
  • 15. Substances Homogeneous materials that always have the same composition.
  • 16. Examples Pure Sugar Pure Salt
  • 17. Substances can be divided into two categories:
    • Elements – substances composed of only one kind of atom
      • Examples – sulfur, oxygen, hydrogen, copper, and gold
    • Compounds – substances composed of more than one kind of atom
      • Example – Water, H 2 O (Atoms in a compound are always in definite proportion, like 2 hydrogen to 1 oxygen in water.)
  • 18. Organic vs. Inorganic
    • Organic compounds mean that carbon is contained.
    • Inorganic compounds mean that no carbon is contained
    • There are a few exceptions…
  • 19. Physical and Chemical Changes A quick review…
  • 20. Physical Changes
    • A physical change occurs when a substance is subjected to some condition, and the substance remains.
    • Examples:
      • Pounding copper sheets
      • Cutting wood
      • Tearing Paper
      • Dissolving sugar in water
  • 21. Chemical Changes
    • Whenever a substance undergoes a change so that one or more new substance with different characteristics is formed, a chemical change (or chemical reaction) has taken place.
      • *A hint… If a precipitate, gas, color change, or energy change occurs, a chemical change has taken place. (There are some exceptions.)
      • A precipitate is a solid substance that forms from a solution.
  • 22. Physical and Chemical Properties A quick review…
  • 23. Physical Properties
    • A physical property is a description of the behavior of a substance undergoing a physical change.
      • Extensive properties – depend on the amount of matter present
        • Mass, length, and volume
      • Intensive properties – do not depend on the amount of matter present
        • Density, malleability, ductility, conductivity, color, melting point, and boiling point
  • 24. Chemical Properties
    • A chemical property describes the reaction of a substance with other materials such as air, water, acid, or a reaction within the substances itself.
      • Example: Iron and water  rust
  • 25. Energy Transfer
  • 26.
    • The most common form of energy change involves heat.
    • Heat is the energy transferred as a result of a temperature difference and is represented by the letter, q.
    • Two ways that heat can be transferred:
      • Contact
        • - Energy will transfer from matter with a higher temp to an object with a lower temp until the objects are equal in temp
      • Work
        • - Surroundings can do work on a system
  • 27. Quantitative measurements of energy changes are expressed in joules, J. Calories are used to measure energy changes, too. 1 calorie = 4.184 joules
  • 28. Energy and Chemical Changes
    • Chemical changes are always accompanied by a change in energy.
    • Two types of reactions:
      • Endothermic – energy is absorbed
        • These reactions get cold because they release no heat.
      • Exothermic – energy is released
        • These reactions get hot because they are giving off energy.
  • 29. Activation Energy
    • Both of these reactions require a certain amount of energy to get started called activation energy .
    • Example
    • Striking a match – friction is the activation energy, causing an exothermic reaction.
  • 30. Measuring Energy Changes
    • A calorimeter is a device used to measure the energy given off or absorbed during a chemical/physical change.
    • To change the temp of a substance, heat must be added or removed.
  • 31.
    • Some substances require little heat, while others require a lot for the same temp change
    • Example
      • 1 gram of liquid water needs 4.184 J of heat to raise its temp 1 ˚C
      • 1 gram of aluminum needs only 0.902 J to raise its temp 1 ˚C
  • 32. Specific Heat
    • The heat needed to raise 1 gram of a substance by one degree Celsius is called its specific heat (C p ).
    • Every substance has its own C p
      • Example
      • The heat required to raise the temp of 1g of water 1 ˚C is 4.184 J. The C p of water is 4.184 J/g ·C˚ (joule per gram Celsius degree).
  • 33. Specific Heats (con’t)
    • Specific heats can be used to find the change in temp of a specific mass of a substance.
    • The Law of Conservation of Energy states that energy is always conserved.
    • So, heat lost by one quantity of matter is gained by another through a energy transfer.
  • 34. q = m( Δ T)(C p )
    • heat gained/lost = mass in grams x change in temp x specific heat
    • Δ T = change in temperature
    • - T final – T initial  when heat is gained
    • - T initial – T final  when heat is lost
  • 35. Problem
    • How much heat is lost when a solid aluminum ingot with a mass of 4110 g cools from 660.0 ˚C to 25˚C?
    • Given:
    • m = 4110 g
    • Δ T = T initial – T final = 660.0 ˚C - 25˚C = 635˚C
    • C p = 0.903 J/g ·C˚
    • Unknown – q = ?
  • 36.
    • Equation
    • q = m( Δ T)(C p )
    • q = (4110 g)(635˚C)( 0.903 J )
    • g ·C˚
    • q = 2,400,000J
  • 37.
    • Suppose a piece of iron with a mass of 21.5g at a temp of 100.0 ˚C is dropped into an insulated container of water. The mass of the water is 132g and its temp before adding the iron is 20.0˚C. What will be the final temp of the system?
    • Given:
    • m iron = 21.5g T initial = 100.0 ˚C
    • m water = 132g T final = 20.0˚C
    • Unknown: T final = ?
    • Equation: q = m( Δ T)(C p )
  • 38. Step 1
    • Heat lost by the iron
    • q = m( Δ T)(C p )
    • q = (21.5g)(100.0˚C - 20.0˚C)( 0.449J )
    • g ·C˚
  • 39. Step 2
    • Heat gained by water
    • q = m( Δ T)(C p )
    • q = (132g)(Tf - 20.0˚C)( 4.184J )
    • g ·C˚
  • 40. Step 3
    • Heat gained must equal heat lost
    • (132g)(Tf - 20.0˚C)( 4.184J ) = (21.5g)(100.0˚C - 20.0˚C)( 0.449J )
    • g ·C g ·C˚