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2 Unit 2: Classification of Matter

3 Unit 2 “Chemical change has always been part of the Universe, even before human beings evolved. Indeed, scientists believe that life began on Earth as a result of complex chemicals reproducing themselves over billions of years. Chemistry is a physical science; it lies between the biological sciences helping to explain many of life’s processes, and the laws of physics, which include matter and energy. Chemical processes are constantly occurring within us – when our bodies move, a series of chemical reactions takes place to give the muscles the energy that is taken in from food. Many species of the animal world make use of chemistry to defend themselves, to kill their prey, and to build fragile structures that have incredible strength. Modern methods of chemical analysis have led to greater understanding of the chemistry of nature, so that it is possible to identify those chemical compounds that produce the color, taste, and smell of a flower or a fruit.” Eyewitness Science “Chemistry”, Dr. Ann Newmark, DK Publishing, Inc., 1993, pg 8

4 Classification of Matter MATTER (gas. Liquid, solid, plasma) PURE SUBSTANCES MIXTURES HETEROGENEOUS MIXTURE HOMOGENEOUS MIXTURES ELEMENTSCOMPOUNDS Separated by physical means into Separated by chemical means into Kotz & Treichel, Chemistry & Chemical Reactivity, 3 rd Edition, 1996, page 31

5 Matter Substance Definite composition (homogeneous) Substance Definite composition (homogeneous) Element (Examples: iron, sulfur, carbon, hydrogen, oxygen, silver) Element (Examples: iron, sulfur, carbon, hydrogen, oxygen, silver) Mixture of Substances Variable composition Mixture of Substances Variable composition Compound (Examples: water. iron (II) sulfide, methane, Aluminum silicate) Compound (Examples: water. iron (II) sulfide, methane, Aluminum silicate) Homogeneous mixture Uniform throughout, also called a solution (Examples: air, tap water, gold alloy) Homogeneous mixture Uniform throughout, also called a solution (Examples: air, tap water, gold alloy) Heterogeneous mixture Nonuniform distinct phases (Examples: soup, concrete, granite) Heterogeneous mixture Nonuniform distinct phases (Examples: soup, concrete, granite) Chemically separable Physically separable

6 The Organization of Matter MATTER PURE SUBSTANCES HETEROGENEOUS MIXTURE HOMOGENEOUS MIXTURES ELEMENTSCOMPOUNDS Physical methods Chemical methods Zumdahl, Zumdahl, DeCoste, World of Chemistry  2002, page 41

7 1) Chemistry: The study of matter, the changes matter goes through, and the associated energy changes. 2) Matter: Anything that has mass and takes up space (volume)

8 Pure Substance A sample of matter in which all parts have the same properties Mixture 2 or more substances that are physically combined – individual properties are retained

9 Element A substance that can’t be broken down into any other substance by ordinary chemical change Compound A substance made of 2 or more elements chemically combined

10 Homogeneous Consistent properties throughout Heterogeneous Uneven, inconsistent distribution of particles

11 8) Mixtures vs Elements & Cmpds a)Mixtures retain properties of constituents b)Composition of a mixture can vary c)Mixtures can be homogeneous or heterogeneous

12 10) Ways to Make Mixtures a)Element + 1 or more elements b)Cmpd + 1 or more elements c)Cmpd + 1 or more cmpds

13 Gold 24 karat gold 18 karat gold 14 karat gold Gold Copper Silver 18 / 24 atoms Au 24 / 24 atoms Au 14 / 24 atoms Au

14 Elements, Compounds, and Mixtures 12) If you have H 2 and O 2 in a container, do you have water? (a) an element (hydrogen) (b) a compound (water) (c) a mixture (hydrogen and oxygen) (d) a mixture (hydrogen and oxygen) Dorin, Demmin, Gabel, Chemistry The Study of Matter, 3 rd Edition, 1990, page 68 hydrogen atoms hydrogen atoms oxygen atoms = Oxygen atom= hydrogen atom

15 Matter Pure Substance Definite composition (homogeneous) Pure Substance Definite composition (homogeneous) Element (Examples: iron, sulfur, carbon, hydrogen, oxygen, silver) Element (Examples: iron, sulfur, carbon, hydrogen, oxygen, silver) Mixture of Substances Variable composition Mixture of Substances Variable composition Compound (Examples: water. iron (II) sulfide, methane, Aluminum silicate) Compound (Examples: water. iron (II) sulfide, methane, Aluminum silicate) Homogeneous mixture Uniform throughout, also called a solution (Examples: air, tap water, gold alloy) Homogeneous mixture Uniform throughout, also called a solution (Examples: air, tap water, gold alloy) Heterogeneous mixture Nonuniform distinct phases (Examples: soup, concrete, granite) Heterogeneous mixture Nonuniform distinct phases (Examples: soup, concrete, granite) Chemically separable Physically separable How would you categorize elements and compounds?

16 During a “physical change” a substance changes some physical property… During a “physical change” a substance changes some physical property… H2OH2O

17 …but it is still the same material with the same chemical composition. H2OH2O gas solid liquid

18 Physical and Chemical Properties Examples of Physical Properties – properties that can be observed without changing the substance Boiling point Color SlipperinessElectrical conductivity Melting point TasteOdorDissolves in water Shininess (luster) SoftnessDuctilityViscosity (resistance to flow) Volatility HardnessMalleabilityDensity (mass / volume ratio) Examples of Chemical Properties – the ability of a substance to undergo a chemical reaction Burns in air Reacts with certain acidsDecomposes when heated Explodes Reacts with certain metalsReacts with certain nonmetals Tarnishes Reacts with waterIs toxic Ralph A. Burns, Fundamentals of Chemistry 1999, page 23 Chemical properties can ONLY be observed during a chemical reaction!

19 The formation of a mixture The formation of a compound Chemical Change Physical Change

20 Physical & Chemical Changes Limestone, CaCO 3 crushing PHYSICAL CHANGE Crushed limestone, CaCO 3 heating CHEMICAL CHANGEPyrex CO 2 CaO Lime and carbon dioxide, CaO + CO 2

21 Pyrex O2O2 H2OH2OPyrex H2O2H2O2 Light hastens the decomposition of hydrogen peroxide, H 2 O 2. The dark bottle in which hydrogen peroxide is usually stored keeps out the light, thus protecting the H 2 O 2 from decomposition. Sunlight energy

22 Density

23 Properties of Matter Pyrex Pyrex Extensive Properties Intensive Properties volume: mass: density: temperature: 100 mL g g/mL 20 o C 15 mL g g/mL 20 o C

24 Solubility – a measure of the amount of solute that can be dissolved in a solvent at a given temperature. Dissolving of Salt in Water NaCl(s) + H 2 O  Na + (aq) + Cl - (aq)

25 Dissolving of NaCl Timberlake, Chemistry 7 th Edition, page 287 HH O Na Cl hydrated ions

26 Density Density is an INTENSIVE INTENSIVE property of matter. - does NOT depend on quantity of matter. - color, melting point, boiling point, odor, density Contrast with EXTENSIVE - depends on quantity of matter. - mass, volume, heat content (calories) Styrofoam Brick

27 StyrofoamBrick ? It appears that the brick is ~40x more dense than the Styrofoam.

28 M M M M V V == D D V V D D Brick Styrofoam Brick

29 Determining the volume of an irregular solid V final = 98.5 cm 3 - V initial = 44.5 cm 3 V fishing sinker = 54.0 cm 3 Before immersion Water 44.5 cm 3 After immersion Fishing sinker 98.5 cm 3 Thread

30 Density D M V ensity ass olume D = M VM V M = D x V V = M DM D

31 Cube Representations 1 m 3 = cm 3 Zumdahl, Zumdahl, DeCoste, World of Chemistry  2002, page 119

32 Consider Equal Volumes The more massive object (the gold cube) has the _________ density. Equal volumes… …but unequal masses aluminum gold GREATER Density = Mass Volume Dorin, Demmin, Gabel, Chemistry The Study of Matter, 3 rd Edition, 1990, page 71

33 Consider Equal Masses Equal masses… …but unequal volumes. The object with the larger volume (aluminum cube) has the density. aluminum gold Dorin, Demmin, Gabel, Chemistry The Study of Matter, 3 rd Edition, 1990, page 71 smaller Christopherson Scales Made in Normal, Illinois USA

34 Consider Equal Masses Equal masses… …but unequal volumes. The object with the larger volume (aluminum cube) has the density. aluminum gold Dorin, Demmin, Gabel, Chemistry The Study of Matter, 3 rd Edition, 1990, page 71 smaller Christopherson Scales

35 Two ways of viewing density Dorin, Demmin, Gabel, Chemistry The Study of Matter, 3 rd Edition, 1990, page 71 Equal volumes… …but unequal masses The more massive object (the gold cube) has the greater density. aluminum gold (A) Equal masses… …but unequal volumes. (B) gold aluminum The object with the larger volume (aluminum cube) has the smaller density.

36 Density of Some Common Substances Density of Some Common Substances Substance Density (g / cm 3 ) Air * Lithium 0.53 Ice Water 1.00 Aluminum 2.70 Iron 7.86 Lead 11.4 Gold 19.3 Density of Some Common Substances Substance Density (g / cm 3 ) Air * Lithium 0.53 Ice Water 1.00 Aluminum 2.70 Iron 7.86 Lead 11.4 Gold 19.3 *at 0 o C and 1 atm pressure

37 Which liquid has the highest density? Coussement, DeSchepper, et al., Brain Strains Power Puzzles  2002, page 16 least dense 1 < 3 < 5 < 2 < 4 most dense

38 Specific Gravity Jaffe, New World of Chemistry, 1955, page water 1.0 ice cork aluminum 2.7 It is a unitless quantity that expresses a ratio of the substance’s density compared to the reference’s density

39 Chapter 13: States of Matter Solid, Liquid, and Gas (13.3) (13.2) (13.1) SolidLiquidGas Zumdahl, Zumdahl, DeCoste, World of Chemistry  2002, page 441 Gas Liquid Solid

40 Average Kinetic Energy and Temperature – they’re the same thing Kinetic energy Fractions of particles Average KE 1 = lower temperature Average KE 2 = higher temperature minimum energy for reaction

41 Hot vs. Cold Tea Kinetic energy Many molecules have an intermediate kinetic energy Few molecules have a very high kinetic energy Low temperature (iced tea) High temperature (hot tea) Percent of molecules ~ ~ ~

42 SOLIDS – Chapter 13, Section 3 (p ) True solids, or crystaline solids, have a crystal lattice structure – a characteristic, geometric arrangement of particles in a solid

43 Amorphous (Glass) Crystalline

44 Amorphous Particle Arrangement Regular, geometric lattice structure Irregular, no specific internal order ShapeCharacteristic crystal structure Irregular (broken glass) Attractive Forces StrongVariable / Weak Melting PointDefinite, specificIndefinite – softens gradually ExamplesQuartzGlass DiamondButter

45 Allotropes of Carbon Graphite

46 Allotropes of Carbon Diamond Zumdahl, Zumdahl, DeCoste, World of Chemistry  2002, page 455

47 Diamond

48 Sodium Chloride Crystal Zumdahl, Zumdahl, DeCoste, World of Chemistry  2002, page 455 = Cl - = Na + crystal lattice structure a characteristic, geometric arrangement of particles in a solid

49 Macromolecules and Allotropes of Carbon GraphiteBuckminsterfullereneDiamond Zumdahl, Zumdahl, DeCoste, World of Chemistry  2002, page 27

50 Allotropes of Carbon C 60 & C 70 “Buckytubes” Buckminsterfullerene “Buckyballs”

51 Fullerenes

52

53 Credit: Baughman et al., Science 297, 787 (2002)

54 Trojan Horse Can use ‘camouflage’ to hide things. Be careful what’s in the Trojan! Buckyballs can hide medicine to treat the human body.

55 Solid H 2 O (s) Ice Zumdahl, Zumdahl, DeCoste, World of Chemistry  2002, page 31

56 Liquid H 2 O (l) Water Zumdahl, Zumdahl, DeCoste, World of Chemistry  2002, page 31 In a liquid molecules are in constant motion there are appreciable intermolecular forces molecules are close together Liquids are almost incompressible Liquids do not fill the container

57 Gas H 2 O (g) Steam Zumdahl, Zumdahl, DeCoste, World of Chemistry  2002, page 31

58 Characteristics of Gases Gases expand to fill any container. –random motion, no attraction Gases are fluids (like liquids). –no attraction Gases have very low densities. –no volume = lots of empty space Courtesy Christy Johannesson

59 Phase Changes

60 Kinetic Theory

61 Kinetic Molecular Theory Postulates Evidence 1. Gases are tiny molecules in mostly empty space. The compressibility of gases. 2. The molecules move in constant, rapid, random, straight-line motion. Gases mix rapidly. 3. The molecules collide elastically with container walls and one another. Gases exert pressure that does not diminish over time. 4. There are no attractive forces between molecules. Gases do not clump. 5. The average kinetic energy of the molecules is proportional to the Kelvin temperature of the sample. Charles’ Law

62 Kinetic Molecular Theory The main assumptions of KMT are… –Gases are made of tiny, individual particles. –The particles move in rapid, random straight-line motion. –Collisions between particles are elastic Courtesy Christy Johannesson

63 Kinetic Molecular Theory Particles in an ideal gas… –have no attractive forces between them. –have no individual volume. Courtesy Christy Johannesson

64 Kinetic Molecular Theory Particles in an ideal gas… –have no attractive forces between them. –have no individual volume. Why do we use a gas that doesn’t really exist? –It simplifies the model and makes it easier to use –Under most normal conditions real gases behave like ideal gases Courtesy Christy Johannesson

65 Real Gases Particles in a REAL gas… –have their own volume –attract each other Gas behavior is most ideal… –at low pressures –at high temperatures –in nonpolar atoms/molecules Courtesy Christy Johannesson

66 8 Elastic vs. Inelastic Collisions 8 v1v1 elastic collision inelastic collision v2v2 v3v3 v4v4

67 Model Gas Behavior All collisions must be elastic Take one step per beat of the metronome Container –Class stands outside tape box Higher temperature –Faster beats of metronome Decreased volume –Divide box in half More Moles –More students are inside box  Mark area of container with tape on ground.  Add only a few molecules of inert gas  Increase temperature  Decrease volume  Add more gas  Effect of diffusion  Effect of effusion (opening size)

68 Average Kinetic Energy and Temperature – they’re the same thing Kinetic energy Fractions of particles Average KE 1 = lower temperature Average KE 2 = higher temperature minimum energy for reaction

69 Microscopic view of a liquid near its surface The high energy molecules escape the surface. Kinetic energy Fractions of particles Average KE 1 = lower temperature Average KE 2 = higher temperature minimum energy to change phase

70 Evaporation H 2 O(g) molecules (water vapor) H 2 O(l) molecules A dynamic equilibrium can only be achieved in a closed container

71 Liquid/Vapor Dynamic Equilibrium The two key properties we need to describe are EVAPORATIONCONDENSATION EVAPORATION and its opposite CONDENSATION H 2 O (l ) → H 2 O (g) H 2 O (g) → H 2 O (l ) At equilibrium the rate of evaporation is equal to the rate of condensation add energy and break intermolecular attractions EVAPORATION release energy and form intermolecular attractions CONDENSATION

72 Water Molecules in Liquid and Steam

73 Microscopic view of a liquid near its surface The high energy molecules escape the surface. Evaporation can only take place at the surface of the liquid At higher temperatures, more particles can escape intermolecular attractions

74 Behavior of a liquid in a closed container A dynamic equilibriumcan only be achieved in a closed container

75 To evaporate, molecules must have sufficient energy to break IM forces. Molecules at the surface break away and become gas. Only those with enough KE escape. endothermicBreaking IM forces requires energy. The process of evaporation is endothermic. Evaporation is a cooling process. It requires heat. Evaporation

76 Change from gas to liquid Achieves a dynamic equilibrium with vaporization in a closed system. What is a closed system? A closed system means matter can’t go in or out. (put a cork in it) What the heck is a “dynamic equilibrium?” Condensation

77 When first sealed, the molecules gradually escape the surface of the liquid. As the molecules build up above the liquid - some condense back to a liquid. The rate at which the molecules evaporate and condense are equal. Dynamic Equilibrium

78 As time goes by the rate of vaporization remains constant but the rate of condensation increases because there are more molecules to condense. Equilibrium is reached when: Rate of Vaporization = Rate of Condensation Molecules are constantly changing phase “dynamic” The total amount of liquid and vapor remains constant “equilibrium” Dynamic Equilibrium

79 Vapor Pressure more “sticky” less likely to vaporize In general: LOW v.p. not very “sticky” more likely to vaporize In general: HIGH v.p.  measure of the tendency for liquid particles to enter gas phase at a given temp.  a measure of “stickiness” of liquid particles to each other NOT all liquids have same v.p. at same temp.

80 Vapor Pressure of: 40°C = 40°C = 40°C = 58 kPa 17 kPa 8 kPa kPa is Standard Atmospheric Pressure

81 Copyright © 2007 Pearson Benjamin Cummings. All rights reserved. Component Percent composition Nitrogen, N 2 78% Oxygen, O 2 21% Argon, Ar 0.9% Water, H 2 O 0 – 4% (variable) Carbon dioxide, CO % (variable) The Earth’s Atmosphere From Space

82 Pressure KEY UNITS AT SEA LEVEL kPa (kilopascal) 1 atm 760 mm Hg 760 torr 14.7 psi 1013 mbar 14.7 psi Courtesy Christy Johannesson Sea level

83 Formation of a bubble is opposed by the pressure of the atmosphere When the vapor pressure is equal to atmospheric pressure the bubble can expand and the liquid boils Zumdahl, Zumdahl, DeCoste, World of Chemistry  2002, page 452

84 TEMPERATURE ( o C) PRESSURE (kPa) CHLOROFORM ETHANOL WATER Volatile substances evaporate easily (have high v.p.’s). BOILING  when vapor pressure = confining pressure (usually from atmosphere) b.p. = 78 o C b.p. = 100 o C atmospheric pressure is kPa

85 Vapor Pressure o C78.4 o C100 o C chloroform ethyl alcohol water Pressure (KPa) Temperature ( o C) 101.3

86 Boiling vs. Evaporation Boiling point: temperature at which vapor pressure = atmospheric pressure Revolutionary process - fast AIR PRESSURE 90 kPa VAPOR PRESSURE 90 kPa Normal Boiling point: temperature at which vapor pressure = standard atmospheric pressure AIR PRESSURE kPa VAPOR PRESSURE kPa

87 Evaporation vs. Boiling Boiling: Change from liquid to gas at the boiling point temperature Evaporation: molecules change from liquid to gas phase below boiling point temperature.

88 If water is boiling at 89°C, what is the pressure? Think: if a liquid is boiling, then it MUST have a VAPOR PRESSURE = TO ATMOSPHERIC PRESSURE kPa is Standard Atmospheric Pressure The pressure is approximately 65 kPa

89 Boiling Point of: 80 kPa = 80 kPa = 80 kPa = 48°C kPa is Standard Atmospheric Pressure 73°C 94°C

90 Boiling Point on Mt. Everest Water exerts a vapor pressure of kPa at a temperature of 100 o C. This is defined as its normal boiling point: ‘vapor pressure = atmospheric pressure’ o C78.4 o C100 o C chloroform ethyl alcohol water Temperature ( o C) Pressure (KPa) On top of Mt. Everest 760 mm Hg x kPa = 253 mm Hg kPa =33.7 kPa

91 Why is boiling a cooling process? Which particles would change phase and how would that effect the average KE? Kinetic energy Fractions of particles minimum energy To change phase Average KE = temperature

92 Liquefaction A gas will change from a gas to a liquid under conditions of: High pressure = particles are closer together causing ↑ attractive forces Low temperature = particles move slower allowing more attractive forces to develop

93 Liquefaction A gas will change from a gas to a liquid under conditions of: High pressure = particles are closer together causing ↑ attractive forces Low temperature = particles move slower allowing more attractive forces to develop

94 Heating Curves Energy is added at a constant rate over time Temperature ( o C) Time Melting - PE  Solid - KE  Liquid - KE  Boiling - PE  Gas - KE 

95 Heating Curves Heating Curves (Chapter 17) Temperature Change –change in KE (molecular motion) –depends on the specific heat capacity Specific Heat (C state ) –energy required to raise the temp of 1 gram of a substance by 1°C –water has a very high specific heat capacity Courtesy Christy Johannesson

96 Specific Heats of Some Substances Specific Heat Substance (cal/ g o C)(J/g o C) Water Alcohol Wood Aluminum Sand Iron Copper Silver Gold

97 Heating Curves Phase Change –change in PE (molecular arrangement) –temp remains constant Heat of Fusion (H f ) –energy required to melt 1 gram of a substance at its m.p.

98 Heating Curves Heat of Vaporization (H v ) –energy required to change 1 gram of a substance from liquid to gas –usually larger than  H f …why? Courtesy Christy Johannesson Joule (J ) – The SI unit used to measure the amount of heat absorbed or released during a reaction.

99 Calculating Energy Changes - Heating Curve for Water q = heat m = mass C state = specific heat ΔT = change in temp Temperature ( o C) Time q = m x H f q = m x H v q = m x C water x  q = m x C steam x  q = m x C ice x  C steam = 1.86 J / g °C H f = 334 J / g H v = 2260 J / g C water = 4.18 J / g °C C ice = 2.14 J / g °C

100 112: Calculating Energy Changes - Heating Curve for Water How many Joules of heat are needed to completely melt 100g of ice at 0°C to water at 0°C ? Temperature ( o C) Time q = m x H f H f = 334 J / g q = m H f q = (100g)(334 J/g) q = 33,400 J

101 113: Calculating Energy Changes - Heating Curve for Water How many Joules are needed to completely boil 200g of water at 100°C to steam at 100°C ? Temperature ( o C) Time q = m x H v H v = 2260 J / g q = m H v q = (200g)(2260 J/g) q = 452,000 J

102 114: Calculating Energy Changes - Heating Curve for Water How many Joules are needed to change the temperature of 50g of water from 25°C to 95°C ? Temperature ( o C) Time q = m x C water x  C water = 4.18 J / g °C q = m C water ΔT q = (50g)(4.18J/g °C)(70°C) q = 14,630 J ΔT = 95°C - 25°C = 70°C

103 Cooling Curve for Water – energy is being removed at a constant rate

104 Calorimetry

105 Vaporization is an endothermic process - it requires heat. Energy is required to overcome intermolecular forces Responsible for cool earth Why we sweat Vaporization

106 solid liquid gas Heat added Temperature ( o C) A B C D E Heating Curve for Water LeMay Jr, Beall, Robblee, Brower, Chemistry Connections to Our Changing World, 1996, page 487

107 Energy Changes Accompanying Phase Changes Exothermic – Energy is released Endothermic – Energy is absorbed Solid Liquid Gas Melting Freezing Deposition CondensationVaporization Sublimation Energy of system Endothermic changes Brown, LeMay, Bursten, Chemistry  2000, page 405 Energy of system Exothermic changes

108 solid liquid gas vaporization condensation melting freezing Heat added Temperature ( o C) A B C D E Heating Curve for Water LeMay Jr, Beall, Robblee, Brower, Chemistry Connections to Our Changing World, 1996, page 487 Endothermic Changes Exothermic Changes

109 Caloric Values Food joules/grams calories/gram Calories/gram Protein Fat Carbohydrates Smoot, Smith, Price, Chemistry A Modern Course, 1990, page calories = 1 Calorie "science" "food" 1calories = joules

110 A Coffee Cup Calorimeter Zumdahl, Zumdahl, DeCoste, World of Chemistry  2002, page 302 Thermometer Styrofoam cover Styrofoam cups Stirrer Thermometer Glass stirrer Cork stopper Two Styrofoam ® cups nested together containing reactants in solution

111 A Bomb Calorimeter

112 Molecular Structure of Ice Hydrogen bonding Zumdahl, Zumdahl, DeCoste, World of Chemistry  2002, page 455

113

114 Alloys

115 Solid Brass An alloy is a mixture of metals. Brass = Copper + Zinc Solid brass homogeneous mixture a substitutional alloy Copper Zinc

116 Brass Plated Brass = Copper + Zinc Brass plated heterogeneous mixture Only brass on outside Copper Zinc

117 Hardened Steel Iron Carbon Steel a interstitial alloy

118 Brass a substitutional alloy Copper Zinc Carbon steel an interstitial alloy Carbon Iron

119 Steel Alloys Stainless steel Tungsten hardened steel Vanadium steel We can engineer properties –Add carbon to increase strength –Too much carbon  too brittle and snaps –Too little carbon  too ductile and iron bends Tensile strength Force is added

120 Separation Techniques

121 Reaction Rates Amount of reactant used or product produced per unit of time A catalyst is a substance that increases the rate of a chemical reaction.catalyst Ex., Enzymes, catalytic convertorsEnzymescatalytic convertors An inhibitor is a substance that decreases the rate of a chemical reaction. Ex., Preservatives

122 Elements, Compounds and Mixtures Classification of Matter MATTER (gas. Liquid, solid, plasma) PURE SUBSTANCES MIXTURES HETEROGENEOUS MIXTURE HOMOGENEOUS MIXTURES ELEMENTSCOMPOUNDS Separated by physical means into Separated by chemical means into Kotz & Treichel, Chemistry & Chemical Reactivity, 3 rd Edition, 1996, page 31

123 Classification of Matter uniform properties? fixed composition? chemically decomposable? no yes hetero- geneous mixture solution element compound

124 The Chemical Parts List

125 Diatomic Elements, 1 and 7 H2H2 N2N2 O2O2 F2F2 Cl 2 Br 2 F2F2 The 7 Diatomic Elements

126 Chemical Formulas – shorthand notation for a chemical compound 5 H 2 SO 4 Subscript – tells the number of atoms of the preceding element in one formula unit Coefficient – tells the number of formula units in the entire expression Chemical Formula – tells the number of atoms of each element in one formula unit of the substance Number of atoms of an element in an each formula unit is given by the subscript. Coefficient x Subscript = Number of atoms in an expression

127 What Do Chemical Formulas Represent?

128 Chemical Formulas – shorthand notation for a chemical compound Expression# of formula units # of atoms of each type in a formula unit # of atoms of each type in entire expression Total atoms in the expression 2 NaOH NH 4 OH 5 Ca(OH) 2 Subscripts outside of parentheses apply to everything inside the parentheses Na = 1 O = 1 H = 2 Na = 2 O = 2 H = 4 Ca = 1 O = 2 x 1 = 2 H = 2 x 1 = 2 N = 1 O = 1 H = 4 +1 = 5 N = 1 O = 1 H = 4 +1 = 5 Ca = 5 O = 10 H = 10

129 Methods of Separating Mixtures Filtration Evaporation Distillation Fractional Distillation Crystallization Decanting

130 Filtration separates a liquid from an insoluble solid Zumdahl, Zumdahl, DeCoste, World of Chemistry  2002, page 40 Mixture of solid and liquid Stirring rod Filtrate (liquid component of the mixture) Filter paper traps solid Funnel

131 Paper Chromatography

132 Separation by Chromatography sample mixture a chromatographic column stationary phase selectively absorbs components mobile phase sweeps sample down column detector

133 Separation by Chromatography sample mixture a chromatographic column stationary phase selectively absorbs components mobile phase sweeps sample down column detector

134 Ion chromatogram of orange juice time (minutes) detector response Na + K+K+ Mg 2+ Fe 3+ Ca 2+

135 A Distillation Apparatus used to separate a soluble solid from a liquid when either or both are to be retained liquid with a solid dissolved in it thermometer condenser tube distilling flask pure liquid receiving flask hose connected to cold water faucet Dorin, Demmin, Gabel, Chemistry The Study of Matter, 3 rd Edition, 1990, page 282

136 The solution is boiled and steam is driven off. Zumdahl, Zumdahl, DeCoste, World of Chemistry  2002, page 39

137 Salt remains after all water is boiled off. Zumdahl, Zumdahl, DeCoste, World of Chemistry  2002, page 39

138 No chemical change occurs when salt water is distilled. Zumdahl, Zumdahl, DeCoste, World of Chemistry  2002, page 40 Saltwater solution (homogeneous mixture) Distillation (physical method) Salt Pure water

139 Distillation Copyright © 2007 Pearson Benjamin Cummings. All rights reserved.

140 Setup to heat a solution Zumdahl, Zumdahl, DeCoste, World of Chemistry  2002, page 42 Ring stand Beaker Wire gauze Ring Bunsen burner

141 long spout helps vapors to condense mixture for distillation placed in here Furnace Glass retort A Hero’s Fountain Eyewitness Science “Chemistry”, Dr. Ann Newmark, DK Publishing, Inc., 1993, pg 13

142 Cooling water out Cooling water in Run hose into sink Connect hose to cold water tap

143 Fractional Distillation: separation of two or more miscible liquids based on deiiference in boiling point Used in petroleum industry to separate crude oil into its varoius components

144 Decanting is used to separate a liquid from an insoluble solid when either the liquid, solid or both are to be retained

145 Separation of a sand-saltwater mixture. Zumdahl, Zumdahl, DeCoste, World of Chemistry  2002, page 40

146 Separation of Sand from Salt 1.Gently break up your salt-crusted sand with a plastic spoon. Follow this flowchart to make a complete separation. Salt- crusted sand. Dry sand. Wet sand. Weigh the mixture. Decant clear liquid. Evaporate to dryness. Pour into heat-resistant container. Fill with water. Stir and let settle 1 minute. Weigh sand. Calculate weight of salt. Repeat 3 times? Yes No 2.How does this flow chart insure a complete separation?


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