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Unit 2


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Unit 2

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