Chemistry. What s the point?

Transcription

1 Chemistry What s the point?

2 General Chemistry What is chemistry? The study of matter and its properties, including ways in which energy changes matter.

3 Why Study Chemistry? Uses in everyday life Understand how batteries work, cars burn gas, bleach whitens, drugs cure diseases Understand the dangers in mixing bleach and ammonia, how poisons can be harmful Understand why yeast makes bread rise, why baking soda is used in cookie dough

4 Why Study Chemistry? Careers that need chemistry Chemistry Pharmacy Biology Psychology Nursing Engineering Dietetics Meteorology

5 Why Study Chemistry? Careers that you would not expect to need chemistry, but do Ceramics Environmentalist Painting Architect Photography Auto Technician In general, to understand the world around us, we need to understand some of the basics of chemistry.

6 Units of Study Matter and Energy Gas Laws and KMT Stoichiometry Atomic Structure Periodic Table Chemical Bonding Reactions Solutions Kinetics Acids and Bases Oxidation Reduction Reactions Organic Chemistry Nuclear Chemistry

7 Units of Study NYS Reference Tables for Physical Setting/CHEMISTRY Has lots of information to help This is YOUR COPY Write your name on it and keep it handy and safely in your binder

8 Unit 1 Matter and Energy Unit 1 Matter and Energy Types of Matter Separation of Mixtures Measurement of Heat Energy Temperature Scales Physical and Chemical Changes Phases of Matter

9 Unit 1 Matter and Energy Classification of Matter What is Matter? Matter is any substance that has mass and takes up volume How is mass measured? What are the units? How is volume measured? What are the units?

10 Matter Substances A Substance is a chemically pure material Made up on only one thing Examples Distilled Water 24k Gold Aluminum foil

11 Matter Element An element is a substance composed of only one type of atom Examples Carbon, Hydrogen, oxygen Elements can NOT be decomposed, or broken down further, by chemical means It is the smallest unit of matter

12 Elements All atoms of the same element have the same number of protons The number of protons = the atomic number Symbolized by

13 Elements Several atoms of the same element Note that the circles are all touching

14 Elements Diatomic elements Note that the circles are touching in pairs Br I N Cl H O F

15 Elements Br I N Cl H O F Br 2 I 2 N 2 Cl 2 H 2 O 2 F 2

16 Matter Compound A Compound is a substance composed of more than one type of atom Examples Water, Carbon Dioxide, Sodium Chloride Compounds CAN be decomposed, or broken down further, by chemical means

17 Compounds Compounds have 2 or more elements Note that the circles are touching or or

18 Compounds Note that the circles are touching. Each combination is a different compound containing the same elements, in different ratios.

19 Matter Mixtures Two types of Mixtures Heterogeneous Homogenous

20 Mixtures Heterogeneous Mixture A Heterogeneous mixture contains two or more substances and is non-uniform The different parts can be seen Examples Sand and water Chicken soup Tossed salad

21 Mixtures

22 Mixtures Homogenous Mixture A Homogenous mixture contains two or more substances has a uniform appearance The different parts can NOT be seen Most common homogenous mixtures are Solutions A Solution is a homogenous mixture of 2 or more substances with a uniform appearance.

23 Mixtures Homogenous Mixture Examples: Salt water Milk, especially homogenized milk Each combination is a different mixture with the same elements in different ratios, but not bonded (not touching).

24 Mixtures

25 Mixtures Homogenous Mixture The ratios between elements may vary in a mixture. An Alloy is a homogenous mixture of 2 or more metals Bronze can be either 90% Cu / 10% Sn, or 95% Cu/5% Sn Either mixture is classified as bronze Steel is another example.

26 Mixtures or

27 Mixtures Separation of Mixtures Heterogeneous and homogenous mixtures are separated differently Whole branch of science devoted to separation chemistry is called Chromatography.

28 Mixtures Separation of Heterogeneous Mixtures Uses physical means of separation Filtration uses a filter to collect the solids and separate it from the filtrate. Decanting is the technique of pouring off a top liquid layer from a bottom layer that has settled down.

29 Mixtures Separation of Homogenous Mixtures Uses the differences in the chemical make up of the individual substances. Distillation is the technique of separation based on differences in boiling point between 2 liquids Water boils at 100 C, ethanol boils at 80 C

30 Mixtures

31 Mixtures Separation of Homogenous Mixtures Chromatography is separation based on differences in solubility and molecular mobility Many different types Paper Thin Layer Chromatography (TLC) Liquid Chromatography (LC) Gas Chromatography (GC) High Performance Liquid Chromatography (HPLC)

32 Chemical and Physical Properties Chemical and physical properties refer to the characteristics of substances. Some are easily observable, other can only be seen during a chemical reaction.

33 Chemical and Physical Properties A Physical Property is an observation that can be determined by the senses and without changing the original substance.

34 Chemical and Physical Properties Physical properties include: Color Odor Melting point Boiling point Freezing point Malleability Shape Density Hardness Taste Magnetism Conductivity

35 Chemical and Physical Properties A Physical Change is a change in the appearance of a substance. There is no new substance made Examples: Ice melting to water or water freezing Wood cut in half Candle wax melting

36 Chemical and Physical Properties A Chemical Property is a property that changes during a chemical reaction. It is a fundamental change in the substance itself.

37 Chemical and Physical Properties Chemical properties include: Reactivity with other chemicals Toxicity Oxidation number Flammability Chemical stability Energy required for formation of a compound Heat of combustion Types of chemical bonds that will form

38 Chemical and Physical Changes A Chemical Change is a change in the substance itself. Two or more things combine and form something new Examples: Iron rusting into iron oxide Alka-Seltzer tablet dissolving and releasing CO 2 Wood burning and forming carbon and CO 2

39 Energy Energy Energy is the ability to do work or create heat. It is measured in Joules. It can also be measured as calories, BTUs, faradays

40 Energy Forms of Energy Heat Light Electrical Mechanical Motion These forms can be converted from one to another.

41 Energy Battery Chemical to electrical energy Generator Mechanical to electrical energy Microphone Sound to mechanical to electrical energy

42 Law of Conservation Law of Conservation of Energy Energy can not be created nor destroyed in a chemical reaction. It can be converted from one form to another. Chemical to heat Chemical to electrical

43 Law of Conservation Law of Conservation of Matter Matter can not be created nor destroyed in a chemical reaction. Forms of matter can change, but the amount remains constant.

44 Energy Energy is either given off or absorbed. Energy is a by product of the reaction. Energy is needed to make the reaction happen.

45 Energy Energy that is released is Exothermic. Exo - out An exothermic reaction gives off energy. The temperature of the surroundings will increase. Mixing NaOH and water. In extreme forms, an explosion.

46 Energy Energy that is absorbed is Endothermic. Endo - in An endothermic reaction absorbs energy. The temperature of the surroundings will decrease. Mixing acetonitrile and water. Sports ice packs.

47 Energy Measurement of Energy In order to know if there is a change in energy, some sort of measurement has to be made. Heat energy is measured in joules (J). Table D James P. Joules

48 Energy If energy can be released or absorbed, and it can be measured, then there has to be a calculation. All equations are on Table T. Look for ones to do with heat. q = mcδt q = mh f q = mh v

49 Calculate Energy q = mcδt where q is heat m is mass, in g C is specific heat capacity ΔT is the change in temperature, in C, or T 1 T 2

50 Calculate Energy C is specific heat capacity Specific heat capacity is the amount of heat required to raise 1 g of a substance in temperature by 1 C. Table B for Water

51 Calculate Energy C water = 4.18 J/g per C Each liquid, or solid, has a different specific heat capacity Al is 0.91 J/g C, Cast Iron is 0.46 J/g C Alcohol is 2.4 J/g C

52 Calculate Energy How many joules of heat are absorbed when 30.0 g of water are heated from 30.0 C to 35.0 C?

53 Calculate Energy q = mcδt q = (g water)(spec. heat.cap)(change in temp) q = (30.0 g)(4.18 J/g)(35.0⁰ 30.0⁰) q = 627 J

54 Calculate Energy How many joules of heat are released when 5.0 g of water are cooled from 70.0 C to 50.0 C?

55 Calculate Energy q = mcδt q = (g water)(spec. heat.cap)(change in temp) q = (5.0 g)(4.18 J/g)(70.0⁰ 50.0⁰) q = 418 J

56 Calculate Energy If 10 J of heat are added to 5g of water at 55⁰C, what is the new temperature of the water?

57 Calculate Energy q = mcδt q = (g water)(spec. heat.cap)(change in temp) 10 J = (5g)(4.18 J/g)(ΔT) ΔT = 10/(5)(4.18) = 0.478⁰C = 0.5 ⁰C

58 Calculate Energy ΔT = 0.5 ⁰C is the CHANGE in temperature. But does the temperature go up or down? Original questions asked: If 10 J of heat are added to 5g of water at 55⁰C, what is the new temperature of the water?

59 Calculate Energy Adding heat causes the temperature to go up. ΔT = 55⁰C ⁰C = 55.5 ⁰C

60 Calculate Energy Units for the q= calculation are in Joules. If units are given in kj, then it must be converted to Joules BEFORE the math is done. So how many J are there in 1kJ? 1000

61 Energy Temperature Temperature is a measure of heat flow. Heat flows spontaneously from hot areas or objects to cooler areas or objects.

62 Temperature Temperature is a measure of average kinetic energy. Kinetic Energy is the energy of motion. Something is moving, like air molecules, or water molecules. As motion changes, kinetic energy changes.

63 Kinetic Energy Objects at high temperatures have particles that are moving quickly.

64 Kinetic Energy Objects at low temperatures have particles that are moving slowly.

65 Kinetic Energy Objects at the same temperatures have particles that are moving at the same average speed No change in temperature.

66 Temperature Temperature is a measure of average kinetic energy. Potential Energy is stored energy. Usually in the form of chemical energy.

67 Temperature Kinetic energy is measured by temperature. Temperature must be on a fixed scale. To have a scale, there must be at least 2 points. Two easiest points to get were water freezing and boiling.

68 Temperature Temperature measured in 2 scales. Celsius, C Kelvin, K (notice there is no sign) Both are the same size, or increments, but scales start at different points.

69 Temperature, C Celsius temperature range Water freezes or melts at 0 C Water boils or condenses at 100 C

70 Temperature, K Kelvin temperature range Water freezes or melts at 273K Water boils or condenses at 373K Adds another temperature Absolute zero, 0 K Theoretically the temperature at which all atomic particle movement stops.

71 Temperature Kelvin is the SI unit of temperature, not C Conversion is K = C Reference Tables Table T and Table A

72 Temperature Standard Temperature is 273 K or 0 C, at 1 Atmosphere Pressure STP is 273 K, 1 atm Room temperature is 20 C, which is what value in K? 293K

73 Unit 1 Matter and Energy Unit 1 Matter and Energy Types of Matter Separation of Mixtures Measurement of Heat Energy Temperature Scales Physical and Chemical Changes Phases of Matter

74 Phases of Matter Phases of Matter Gas Liquid Solid Phases change from one to another. Each phase change absorbs or releases energy in a specific manner.

75 Phases Changes Phases change from one to another. Gas Liquid Solid

76 Phases Changes Heating Curves During a heating curve: Energy is absorbed during melting and boiling. These changes are endothermic.

77 Heating Curves During a heating curve Average kinetic energy remains constant. Potential energy increases.

78 Temperature C Heating Curves Boiling Point Gas Liquid and Gas Melting Point Liquid Solid and Liquid Solid Time

79 Temperature C Heating Curves Time

80 Heating Curves All phase changes are endothermic They absorb energy. Melting is also called fusion. Boiling is also called vaporization.

81 Heating Curves AB: heating a solid, 1 phase, kinetic energy (KE) increases. BC: melting a solid, 2 phases, KE constant, potential energy (PE) increases. CD: heating a liquid, 1 phase, KE increases.

82 Heating Curves DE: boiling a liquid, 2 phases, KE constant, PE increases. EF: heating a gas, 1 phase, KE increases. Matter/HeatingCurve.htm

83 Phases Changes Cooling Curves During a cooling curve: Energy is released during condensation and freezing. These changes are exothermic.

84 Cooling Curves During a cooling curve Average kinetic energy remains constant. Potential energy decreases.

85 Temperature C Cooling Curves Gas Liquid and Gas Liquid Condensation Point Freezing Point Solid and Liquid Solid Time

86 Temperature C Cooling Curves Time

87 Cooling Curves All phase changes are exothermic. They release energy. Freezing is also called solidification.

88 Cooling Curves AB: cooling a gas, 1 phase, KE decreases. BC: condensing a gas, 2 phases, KE constant, PE decreases. CD: cooling a liquid, 1 phase, KE decreases.

89 Cooling Curves DE: solidifying a liquid, 2 phases, KE constant, PE decreases. EF: cooling a solid, 1 phase, KE decreases.

90 Phase Changes During phase changes, there is a point where 2 phases coexist. Equilibrium is a state of balance between phases. Melting/freezing at 0 C Condensation/vaporization at 100 C

91 Phase Changes Some substances can go from a solid directly to a gas. Dry ice (CO 2 ), moth balls. Some go from a gas directly to a solid. Water vapor to ice crystals.

92 Phase Changes Sublimation is the process by which a solid goes directly to a gas. Deposition is the process by which a gas goes directly to a solid.

93 Characteristics of Liquids Heat of vaporization, H v is the energy required to boil 1g of water J/g Table B

94 Characteristics of Solids Heat of fusion, or H f, is the energy required to melt 1g of ice. 334 J/g Table B

95 Heat of Vaporization/Fusion q = mh f q = mh v How much energy is required to boil 10 g of water?

96 Heat of Vaporization/Fusion q = mh v q = (10 g)(2260 J/g) = 22,600J

97 Heat of Vaporization/Fusion If 50 g of ice are melted to water, how much heat must be added to the system?

98 Heat of Vaporization/Fusion q = mh f q = (50 g)(334 J/g) = 16,700 J

99 Heat of Vaporization/Fusion 1. What is the mass of water if 25 kj are consumed when the water is boiled? 2. An experiment shows that it requires 20,000 J to melt 32 g of a substance. What is the H v of that substance?

100 Phases of Matter Phases of Matter Phases change from one to another. Each phase change absorbs or releases energy in a specific manner. Solids Liquids Gases (details next unit)

101 Characteristics of Solids Solids Solids have and maintain its own shape and volume. Have a crystalline structure with closely-packed particles in a regular geometric pattern.

102 Characteristics of Solids Supercooled liquids, like glass and Silly-Putty, have definite volume and shape, but lack crystal structure. Referred to as Non-Newtonian fluids.

103 Characteristics of Solids If the temperature is raised, the average kinetic energy increases. As KE, the molecules move by vibrating.

104 Characteristics of Solids Heat of fusion, or H f, is the energy required to melt 1g of ice. 334 J/g Table B

105 Characteristics of Solids High heats of fusion or vaporization indicate strong intermolecular forces. Water has very strong intermolecular forces.

106 Characteristics of Liquids Liquids Liquids take the shape of its container. Keeps its own volume. Not easily compressed.

107 Characteristics of Liquids Evaporation when a liquid changes to a gas at a temperature below its boiling point.

108 Characteristics of Liquids Evaporation happens because some molecules are moving faster than others. If these molecules are near the surface, they escape the attractive forces of the other molecules in the liquid.

109 Characteristics of Liquids In a closed container, the vapor molecules exert pressure on the sides of the container and liquid.

110 Characteristics of Liquids Vapor Pressure is the particle pressure of a vapor at the surface of its parent liquid when in equilibrium. Table H

111 Characteristics of Liquids Equilibrium is a balanced condition resulting from two opposing reactions. Here evaporation and condensation are happening at an equal rate.

112 Characteristics of Liquids If the temperature is raised, the average kinetic energy increases. As KE, so do the number of fast moving molecules that can escape the liquid attraction.

113 Characteristics of Liquids This increases the vapor pressure. As temperature increases, vapor pressure increases.

114 Characteristics of Liquids Vapor pressure is independent of volume. Vapor pressure depends on temperature and pressure.

115 Characteristics of Liquids Which sample of water has the greatest vapor pressure? ml at 20 C ml at 35 C ml at 25 C

116 Temperature C Characteristics of Liquids Boiling point is the temperature at which a liquid becomes a gas. Boiling Point Liquid and Gas Gas Melting Point Solid and Liquid Liquid Solid Time

117 Characteristics of Liquids Table H shows how vapor pressure increases with temperature. Go to the boiling point of water at 100 C. Note that the vapor pressure is equal to standard or atmospheric pressure of kpa.

118 Characteristics of Liquids Any liquid will boil when its vapor pressure is equal to the pressure on its surface. Under lower atmospheric pressures, water will boil at temperatures below 100 C.

119 Characteristics of Liquids Leadville, CO is the highest city in the continental US at 10,430 feet, or 3179 m. The usual air pressure there is about 70 kpa. At what temperature does water boil? Table H about 90 C

120 Characteristics of Liquids Normal boiling point is the temperature at which the vapor pressure of a liquid is equal to standard atmospheric pressure.

121 Characteristics of Liquids At the exact boiling point, there is an equilibrium between vaporization and condensation.

122 Characteristics of Liquids Water is boiling at a pressure of 30 kpa. What is the temperature of the water? 70 C

123 Characteristics of Liquids Water will boil at 15 C when the pressure on its surface is at what kpa? 2 kpa

124 Characteristics of Liquids Intermolecular Forces strong, then BP, VP Intermolecular Forces weak, then BP, VP List from highest (1) to lowest (4) Boiling Point Vapor Pressure Intermolecular Forces

125 Characteristics of Liquids Intermolecular Forces strong, then BP, VP Intermolecular Forces weak, then BP, VP List from highest (1) to lowest (4) Boiling Point Vapor Pressure Intermolecular Forces 1. Ethanoic Acid (Acetic Acid) 2. Water 3. Ethanol 4. Propanone (Acetone)

126 Characteristics of Liquids Intermolecular Forces strong, then BP, VP Intermolecular Forces weak, then BP, VP List from highest (1) to lowest (4) Boiling Point Vapor Pressure Intermolecular Forces 1. Ethanoic Acid (Acetic Acid) 2. Water 2. Ethanol 3. Ethanol 3. Water 4. Propanone (Acetone) 1. Propanone (Acetone) 4. Ethanoic Acid (Acetic Acid)

127 Characteristics of Liquids Intermolecular Forces strong, then BP, VP Intermolecular Forces weak, then BP, VP List from highest (1) to lowest (4) Boiling Point Vapor Pressure Intermolecular Forces 1. Ethanoic Acid (Acetic Acid) 1. Propanone (Acetone) 2. Water 2. Ethanol 2. Water 3. Ethanol 3. Water 3. Ethanol 4. Propanone (Acetone) 4. Ethanoic Acid (Acetic Acid) 1. Ethanoic Acid (Acetic Acid) 4. Propanone (Acetone)

128 Phases of Matter Phases of Matter Phases change from one to another. Each phase change absorbs or releases energy in a specific manner. Liquids and Vapor Pressure Solids Gases and the Ideal Gas Law