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# Chemistry Unit 1 PPT 2

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Chemistry Unit 1 PPT 2 Notes

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### Chemistry Unit 1 PPT 2

1. 1. Chemistry Toombs County High School
2. 2. 3.1  On January 4, 2004, the Mars Exploration Rover Spirit landed on Mars. Each day of its mission, Spirit recorded measurements for analysis. In the chemistry laboratory, you must strive for accuracy and precision in your measurements.
3. 3. Using and Expressing Measurements How do measurements relate to science?
4. 4. A measurement is a quantity that has both a number and a unit.  Measurements are fundamental to the experimental sciences. For that reason, it is important to be able to make measurements and to decide whether a measurement is correct.
5. 5.  In scientific notation, a given number is written as the product of two numbers: a coefficient and 10 raised to a power.  The number of stars in a galaxy is an example of an estimate that should be expressed in scientific notation.
6. 6. Accuracy, Precision, and Error How do you evaluate accuracy and precision?
7. 7. Accuracy and Precision ▪Accuracy is a measure of how close a measurement comes to the actual or true value of whatever is measured. ▪Precision is a measure of how close a series of measurements are to one another.
8. 8. To evaluate the accuracy of a measurement, the measured value must be compared to the correct value. To evaluate the precision of a measurement, you must compare the values of two or more repeated measurements.
9. 9. Determining Error ▪ The accepted value is the correct value based on reliable references. ▪ The experimental value is the value measured in the lab. ▪ The difference between the experimental value and the accepted value is called the error.
10. 10. ▪ The percent error is the absolute value of the error divided by the accepted value, multiplied by 100%.
11. 11.  Just because a measuring device works, you cannot assume it is accurate. The scale below has not been properly zeroed, so the reading obtained for the person’s weight is inaccurate.
12. 12. Significant Figures in Measurements Why must measurements be reported to the correct number of significant figures?
13. 13. ▪ Suppose you estimate a weight that is between 2.4 lb and 2.5 lb to be 2.46 lb. The first two digits (2 and 4) are known. The last digit (6) is an estimate and involves some uncertainty. All three digits convey useful information, however, and are called significant figures. ▪ The significant figures in a measurement include all of the digits that are known, plus a last digit that is estimated.
14. 14. Measurements must always be reported to the correct number of significant figures because calculated answers often depend on the number of significant figures in the values used in the calculation.
15. 15. 3.1
16. 16. 3.1
17. 17.  Significant Figures in Calculations  How does the precision of a calculated answer compare to the precision of the measurements used to obtain it?
18. 18.  In general, a calculated answer cannot be more precise than the least precise measurement from which it was calculated. The calculated value must be rounded to make it consistent with the measurements from which it was calculated.
19. 19. Rounding  To round a number, you must first decide how many significant figures your answer should have. The answer depends on the given measurements and on the mathematical process used to arrive at the answer.
20. 20. Addition and Subtraction  The answer to an addition or subtraction calculation should be rounded to the same number of decimal places (not digits) as the measurement with the least number of decimal places.
21. 21. Multiplication and Division ▪ In calculations involving multiplication and division, you need to round the answer to the same number of significant figures as the measurement with the least number of significant figures. ▪ The position of the decimal point has nothing to do with the rounding process when multiplying and dividing measurements.
22. 22.  In the signs shown here, the distances are listed as numbers with no units attached. Without the units, it is impossible to communicate the measurement to others. When you make a measurement, you must assign the correct units to the numerical value.
23. 23.  Measuring with SI Units  Which five SI base units do chemists commonly use?
24. 24.  All measurements depend on units that serve as reference standards. The standards of measurement used in science are those of the metric system.  The International System of Units (abbreviated SI, after the French name, Le Système International d’Unités) is a revised version of the metric system.
25. 25. The five SI base units commonly used by chemists are the meter, the kilogram, the kelvin, the second, and the mole.
26. 26. Units and Quantities  What metric units are commonly used to measure length, volume, mass, temperature and energy?
27. 27.  Units of Length  In SI, the basic unit of length, or linear measure, is the meter (m). For very large or and very small lengths, it may be more convenient to use a unit of length that has a prefix.
28. 28.  Common metric units of length include the centimeter, meter, and kilometer.
29. 29.  Units of Volume  The SI unit of volume is the amount of space occupied by a cube that is 1 m along each edge. This volume is the cubic meter (m)3 . A more convenient unit of volume for everyday use is the liter, a non-SI unit.  A liter (L) is the volume of a cube that is 10 centimeters (10 cm) along each edge (10 cm × 10 cm × 10 cm = 1000 cm3 = 1 L).
30. 30.  Common metric units of volume include the liter, milliliter, cubic centimeter, and microliter.
31. 31.  The volume of 20 drops of liquid from a medicine dropper is approximately 1 mL.
32. 32. A sugar cube has a volume of 1 cm3 . 1 mL is the same as 1 cm3 .
33. 33.  A gallon of milk has about twice the volume of a 2-L bottle of soda.
34. 34. Units of Mass  The mass of an object is measured in comparison to a standard mass of 1 kilogram (kg), which is the basic SI unit of mass.  A gram (g) is 1/1000 of a kilogram; the mass of 1 cm3 of water at 4°C is 1 g.
35. 35.  Common metric units of mass include kilogram, gram, milligram, and microgram.
36. 36.  Weight is a force that measures the pull on a given mass by gravity.  The astronaut shown on the surface of the moon weighs one sixth of what he weighs on Earth.
37. 37.  Units of Temperature  Temperature is a measure of how hot or cold an object is.  Thermometers are used to measure temperature.
38. 38. ▪ Scientists commonly use two equivalent units of temperature, the degree Celsius and the kelvin.
39. 39. ▪ On the Celsius scale, the freezing point of water is 0°C and the boiling point is 100°C. ▪ On the Kelvin scale, the freezing point of water is 273.15 kelvins (K), and the boiling point is 373.15 K. ▪ The zero point on the Kelvin scale, 0 K, or absolute zero, is equal to −273.15 °C.
40. 40.  Because one degree on the Celsius scale is equivalent to one kelvin on the Kelvin scale, converting from one temperature to another is easy. You simply add or subtract 273, as shown in the following equations.
41. 41. 3.2 ▪ Conversions Between the Celsius and Kelvin Scales
42. 42.  Units of Energy Energy is the capacity to do work or to produce heat. ▪ The joule and the calorie are common units of energy.
43. 43.  The joule (J) is the SI unit of energy.  One calorie (cal) is the quantity of heat that raises the temperature of 1 g of pure water by 1°C.
44. 44.  This house is equipped with solar panels. The solar panels convert the radiant energy from the sun into electrical energy that can be used to heat water and power appliances.
45. 45.  Because each country’s currency compares differently with the U.S. dollar, knowing how to convert currency units correctly is very important. Conversion problems are readily solved by a problem- solving approach called dimensional analysis.
46. 46. Conversion Factors What happens when a measurement is multiplied by a conversion factor?
47. 47. ▪ A conversion factor is a ratio of equivalent measurements. ▪ The ratios 100 cm/1 m and 1 m/100 cm are examples of conversion factors.
48. 48. When a measurement is multiplied by a conversion factor, the numerical value is generally changed, but the actual size of the quantity measured remains the same.
49. 49.  The scale of the micrograph is in nanometers. Using the relationship 109 nm = 1 m, you can write the following conversion factors.
50. 50.  Dimensional Analysis  Why is dimensional analysis useful?
51. 51.  Dimensional analysis is a way to analyze and solve problems using the units, or dimensions, of the measurements.  Dimensional analysis provides you with an alternative approach to problem solving.
52. 52. Converting Between Units  What types of problems are easily solved by using dimensional analysis?
53. 53. Problems in which a measurement with one unit is converted to an equivalent measurement with another unit are easily solved using dimensional analysis.
54. 54.  Multistep Problems  When converting between units, it is often necessary to use more than one conversion factor. Sample problem 3.8 illustrates the use of multiple conversion factors.
55. 55.  Converting Complex Units  Many common measurements are expressed as a ratio of two units. If you use dimensional analysis, converting these complex units is just as easy as converting single units. It will just take multiple steps to arrive at an answer.
56. 56.  If you think that these lily pads float because they are lightweight, you are only partially correct. The ratio of the mass of an object to its volume can be used to determine whether an object floats or sinks in water.
57. 57. Determining Density What determines the density of a substance?
58. 58. Density is the ratio of the mass of an object to its volume.
59. 59.  Each of these 10-g samples has a different volume because the densities vary.
60. 60.  Density is an intensive property that depends only on the composition of a substance, not on the size of the sample.
61. 61. The density of corn oil is less than the density of corn syrup. For that reason, the oil floats on top of the syrup.
62. 62. Density and Temperature How does a change in temperature affect density?
63. 63. Experiments show that the volume of most substances increases as the temperature increases. Meanwhile, the mass remains the same. Thus, the density must change.  The density of a substance generally decreases as its temperature increases.
64. 64.  You could measure the amount of sand in a sand sculpture by counting each grain of sand, but it would be much easier to weigh the sand. You’ll discover how chemists measure the amount of a substance using a unit called a mole, which relates the number of particles to the mass. 10.1
65. 65. Measuring Matter What are three methods for measuring the amount of something? 10.1
66. 66. ▪ You often measure the amount of something by one of three different methods—by count, by mass, and by volume. 10.1
67. 67. What Is a Mole? How is Avogadro’s number related to a mole of any substance? 10.1
68. 68. A mole of any substance contains Avogadro’s number of representative particles, or 6.02 × 1023 representative particles.  The term representative particle refers to the species present in a substance: usually atoms, molecules, or formula units. 10.1
69. 69.  Converting Number of Particles to Moles  One mole (mol) of a substance is 6.02 × 1023 representative particles of that substance and is the SI unit for measuring the amount of a substance.  The number of representative particles in a mole, 6.02 × 1023 , is called Avogadro’s number. 10.1
70. 70. 10.1
71. 71.  Converting Moles to Number of Particles 10.1
72. 72. The Mass of a Mole of an Element How is the atomic mass of an element related to the molar mass of an element? 10.1
73. 73. The atomic mass of an element expressed in grams is the mass of a mole of the element.  The mass of a mole of an element is its molar mass. 10.1
74. 74. One molar mass of carbon, sulfur, mercury, and iron are shown. 10.1
75. 75. 10.1
76. 76. The Mass of a Mole of a Compound  How is the mass of a mole of a compound calculated? 10.1
77. 77.  To calculate the molar mass of a compound, find the number of grams of each element in one mole of the compound. Then add the masses of the elements in the compound. 10.1
78. 78. ▪ Substitute the unit grams for atomic mass units. Thus 1 mol of SO3 has a mass of 80.1 g. 10.1
79. 79. ▪ Molar Masses of Glucose, Water, and Paradichlorobenzene 10.1
80. 80.  How can you guess the number of jelly beans in a jar? You estimate the size of a jelly bean and then estimate the dimensions of the container to obtain its volume. In a similar way, chemists use the relationships between the mole and quantities such as mass, volume, and number of particles to solve chemistry problems. 10.2
81. 81. The Mole–Mass Relationship  How do you convert the mass of a substance to the number of moles of the substance? 10.2
82. 82. ▪ Use the molar mass of an element or compound to convert between the mass of a substance and the moles of a substance. 10.2
83. 83. The Mole–Volume Relationship  What is the volume of a gas at STP? 10.2
84. 84.  Avogadro’s hypothesis states that equal volumes of gases at the same temperature and pressure contain equal numbers of particles. 10.2
85. 85.  The volume of a gas varies with temperature and pressure. Because of these variations, the volume of a gas is usually measured at a standard temperature and pressure.  Standard temperature and pressure (STP) means a temperature of 0°C and a pressure of 101.3 kPa, or 1 atmosphere (atm). 10.2
86. 86. ▪At STP, 1 mol or, 6.02 × 1023 representative particles, of any gas occupies a volume of 22.4 L. ▪The quantity 22.4 L is called the molar volume of a gas. 10.2
87. 87.  Calculating Volume at STP 10.2
88. 88.  Calculating Molar Mass from Density 10.2
89. 89. 10.2
90. 90. 10.2
91. 91. 10.2
92. 92. 10.2
93. 93. Particles Move - KINETIC
94. 94.  The skunk releases its spray! Within seconds you smell that all-too-familiar foul odor. You will discover some general characteristics of gases that help explain how odors travel through the air, even on a windless day. 13.1
95. 95. Kinetic Theory and a Model for Gases What are the three assumptions of the kinetic theory as it applies to gases? 13.1
96. 96. The word kinetic refers to motion.  The energy an object has because of its motion is called kinetic energy.  According to the kinetic theory, all matter consists of tiny particles that are in constant motion. 13.1
97. 97. According to kinetic theory: ▪The particles in a gas are considered to be small, hard spheres with an insignificant volume. ▪The motion of the particles in a gas is rapid, constant, and random. ▪All collisions between particles in a gas are perfectly elastic. 13.1
98. 98. ▪ Particles in a gas are in rapid, constant motion. 13.1
99. 99. ▪ Gas particles travel in straight-line paths. 13.1
100. 100. ▪ The gas fills the container. 13.1
101. 101. Gas Pressure  How does kinetic theory explain gas pressure? 13.1
102. 102.  Gas pressure results from the force exerted by a gas per unit surface area of an object.  An empty space with no particles and no pressure is called a vacuum.  Atmospheric pressure results from the collisions of atoms and molecules in air with objects. 13.1
103. 103. Gas pressure is the result of simultaneous collisions of billions of rapidly moving particles in a gas with an object. 13.1
104. 104.  A barometer is a device that is used to measure atmospheric pressure. 13.1
105. 105.  The SI unit of pressure is the pascal (Pa).  One standard atmosphere (atm) is the pressure required to support 760 mm of mercury in a mercury barometer at 25°C. 13.1
106. 106. Kinetic Energy and Temperature What is the relationship between the temperature in kelvins and the average kinetic energy of particles? 13.1
107. 107. Average Kinetic Energy  The particles in any collection of atoms or molecules at a given temperature have a wide range of kinetic energies. Most of the particles have kinetic energies somewhere in the middle of this range. 13.1
108. 108. 13.1
109. 109.  Absolute zero (0 K, or –273.15°C) is the temperature at which the motion of particles theoretically ceases.  Particles would have no kinetic energy at absolute zero.  Absolute zero has never been produced in the laboratory. 13.1
110. 110. Average Kinetic Energy and Kelvin Temperature The Kelvin temperature of a substance is directly proportional to the average kinetic energy of the particles of the substance. 13.1
111. 111. ▪ In this vacuum chamber, scientists cooled sodium vapor to nearly absolute zero. 13.1
112. 112.  END OF SHOW
113. 113.  In organized soccer, a ball that is properly inflated will rebound faster and travel farther than a ball that is under-inflated. If the pressure is too high, the ball may burst when it is kicked. You will study variables that affect the pressure of a gas.
114. 114.  Compressibility  Why are gases easier to compress than solids or liquids are?
115. 115.  Compressibility is a measure of how much the volume of matter decreases under pressure. When a person collides with an inflated airbag, the compression of the gas absorbs the energy of the impact.
116. 116.  Gases are easily compressed because of the space between the particles in a gas. ▪ The distance between particles in a gas is much greater than the distance between particles in a liquid or solid. ▪ Under pressure, the particles in a gas are forced closer together.
117. 117. ▪ At room temperature, the distance between particles in an enclosed gas is about 10 times the diameter of a particle.
118. 118. Factors Affecting Gas Pressure  What are the three factors that affect gas pressure?
119. 119. The amount of gas, volume, and temperature are factors that affect gas pressure.
120. 120. Four variables are generally used to describe a gas. The variables and their common units are  pressure (P) in kilopascals  volume (V) in liters  temperature (T) in kelvins  the number of moles (n).
121. 121. Amount of Gas  You can use kinetic theory to predict and explain how gases will respond to a change of conditions. If you inflate an air raft, for example, the pressure inside the raft will increase.
122. 122.  Collisions of particles with the inside walls of the raft result in the pressure that is exerted by the enclosed gas. Increasing the number of particles increases the number of collisions, which is why the gas pressure increases.
123. 123.  If the gas pressure increases until it exceeds the strength of an enclosed, rigid container, the container will burst.
124. 124.  Aerosol Spray Paint
125. 125. Volume  You can raise the pressure exerted by a contained gas by reducing its volume. The more a gas is compressed, the greater is the pressure that the gas exerts inside the container.
126. 126.  When the volume of the container is halved, the pressure the gas exerts is doubled.
127. 127. Temperature  An increase in the temperature of an enclosed gas causes an increase in its pressure.  As a gas is heated, the average kinetic energy of the particles in the gas increases. Faster-moving particles strike the walls of their container with more energy.
128. 128.  When the Kelvin temperature of the enclosed gas doubles, the pressure of the enclosed gas doubles.
129. 129.  Hot lava oozes and flows, scorching everything in its path, and occasionally overrunning nearby houses. When the lava cools, it solidifies into rock. The properties of liquids are related to intermolecular interactions. You will learn about some of the properties of liquids. 13.2
130. 130. A Model for Liquids What factors determine the physical properties of a liquid? 13.2
131. 131. ▪ Substances that can flow are referred to as fluids. Both liquids and gases are fluids. 13.2
132. 132. The interplay between the disruptive motions of particles in a liquid and the attractions among the particles determines the physical properties of liquids. 13.2
133. 133. Evaporation What is the relationship between evaporation and kinetic energy? 13.2
134. 134.  The conversion of a liquid to a gas or vapor is called vaporization.  When such a conversion occurs at the surface of a liquid that is not boiling, the process is called evaporation. 13.2
135. 135. 13.2  In an open container, molecules that evaporate can escape from the container.
136. 136.  In a closed container, the molecules cannot escape. They collect as a vapor above the liquid. Some molecules condense back into a liquid. 13.2
137. 137. During evaporation, only those molecules with a certain minimum kinetic energy can escape from the surface of the liquid. 13.2
138. 138. Vapor Pressure  When can a dynamic equilibrium exist between a liquid and its vapor? 13.2
139. 139. Vapor pressure is a measure of the force exerted by a gas above a liquid. 13.2
140. 140.  In a system at constant vapor pressure, a dynamic equilibrium exists between the vapor and the liquid. The system is in equilibrium because the rate of evaporation of liquid equals the rate of condensation of vapor. 13.2
141. 141. Vapor Pressure and Temperature Change ▪ An increase in the temperature of a contained liquid increases the vapor pressure. ▪ The particles in the warmed liquid have increased kinetic energy. As a result, more of the particles will have the minimum kinetic energy necessary to escape the surface of the liquid. 13.2
142. 142. 13.2
143. 143.  Vapor Pressure Measurements ▪ The vapor pressure of a liquid can be determined with a device called a manometer. 13.2
144. 144. ▪ Manometer 13.2
145. 145.  Boiling Point  Under what conditions does boiling occur? 13.2
146. 146. When a liquid is heated to a temperature at which particles throughout the liquid have enough kinetic energy to vaporize, the liquid begins to boil. 13.2
147. 147. The temperature at which the vapor pressure of the liquid is just equal to the external pressure on the liquid is the boiling point (bp). 13.2
148. 148.  Boiling Point and Pressure Changes ▪ Because a liquid boils when its vapor pressure is equal to the external pressure, liquids don’t always boil at the same temperature. ▪ At a lower external pressure, the boiling point decreases. ▪ At a higher external pressure, the boiling point increases. 13.2
149. 149. ▪ Altitude and Boiling Point 13.2
150. 150. 13.2
151. 151.  Normal Boiling Point  Because a liquid can have various boiling points depending on pressure, the normal boiling point is defined as the boiling point of a liquid at a pressure of 101.3 kPa. 13.2
152. 152. 13.2
153. 153.  END OF SHOW
154. 154.  In 1985, scientists discovered a new form of carbon. They called this form of carbon buckminsterfullerene, or buckyball for short. You will learn how the arrangement of particles in solids determines some general properties of solids. 13.3
155. 155. A Model for Solids  How are the structure and properties of solids related? 13.3
156. 156.  The general properties of solids reflect the orderly arrangement of their particles and the fixed locations of their particles. 13.3
157. 157.  The melting point (mp) is the temperature at which a solid changes into a liquid. 13.3
158. 158.  Crystal Structure and Unit Cells  What determines the shape of a crystal? 13.3
159. 159. ▪ In a crystal, the particles are arranged in an orderly, repeating, three-dimensional pattern called a crystal lattice. 13.3
160. 160. Allotropes  Allotropes are two or more different molecular forms of the same element in the same physical state. ▪ Allotropes have different properties because their structures are different. ▪ Only a few elements have allotropes. 13.3
161. 161. ▪ Carbon Allotropes 13.3
162. 162. Non-Crystalline Solids  An amorphous solid lacks an ordered internal structure. ▪ Rubber, plastic, asphalt, and glass are amorphous solids. ▪ A glass is a transparent fusion product of inorganic substances that have cooled to a rigid state without crystallizing. 13.3
163. 163.  END OF SHOW
164. 164. SC .
165. 165.  Familiar weather events can remind you that water exists on Earth as a liquid, a solid, and a vapor. As water cycles through the atmosphere, the oceans, and Earth’s crust, it undergoes repeated changes of state. You will learn what conditions can control the state of a substance. 13.4
166. 166.  Sublimation  When can sublimation occur? 13.4
167. 167. The change of a substance from a solid to a vapor without passing through the liquid state is called sublimation. ▪ Sublimation occurs in solids with vapor pressures that exceed atmospheric pressure at or near room temperature. 13.4
168. 168.  When solid iodine is heated, the crystals sublime, going directly from the solid to the gaseous state. When the vapor cools, it goes directly from the gaseous to the solid state. 13.4
169. 169.  Phase Diagrams  How are the conditions at which phases are in equilibrium represented on a phase diagram? 13.4
170. 170. A phase diagram is a graph that gives the conditions of temperature and pressure at which a substance exists as solid, liquid, and gas (vapor). 13.4
171. 171. The conditions of pressure and temperature at which two phases exist in equilibrium are indicated on a phase diagram by a line separating the phases. 13.4
172. 172. 13.4
173. 173.  The triple point describes the only set of conditions at which all three phases can exist in equilibrium with one another. 13.4
174. 174.  END OF SHOW