1. Entry Task: Sept 5-6 th Block 1 Grab an Entry Task sheet Question: Convert 24001 cm to dam. Provide the name for this acid HC 2 H 3 O 2 You have 5 minutes!

2. Agenda: Ch. 1-2 Review and Clicker

3. Chapter 1 Introduction: Matter and Measurement

4. 1.1 Chemistry: The study of matter and the changes it undergoes.

5. Scientific Method: A systematic approach to solving problems.

6. Matter: Anything that has mass and takes up space.

7. Matter Atoms are the building blocks of matter.

8. Matter Atoms are the building blocks of matter. Each element is made of the same kind of atom.

9. Matter Atoms are the building blocks of matter. Each element is made of the same kind of atom. A compound is made of two or more different kinds of elements.

10. 1.2 States of Matter

11. Classification of Matter

21. Mixtures and Compounds

22. 1.3 Properties and Changes of Matter

23. Properties of Matter Physical Properties: □Can be observed without changing a substance into another substance. Boiling point, density, mass, volume, etc. Chemical Properties: □Can only be observed when a substance is changed into another substance. Flammability, corrosiveness, reactivity with acid, etc.

24. Properties of Matter Intensive Properties: □Independent of the amount of the substance that is present. Density, boiling point, color, etc. Extensive Properties: □Dependent upon the amount of the substance present. Mass, volume, energy, etc.

25. Changes of Matter Physical Changes: □Changes in matter that do not change the composition of a substance. Changes of state, temperature, volume, etc. Chemical Changes: □Changes that result in new substances. Combustion, oxidation, decomposition, etc.

26. Chemical Reactions In the course of a chemical reaction, the reacting substances are converted to new substances.

27. Chemical Reactions

28. Compounds Compounds can be broken down into more elemental particles.

29. Electrolysis of Water

30. Separation of Mixtures

31. Distillation: Separates homogeneous mixture on the basis of differences in boiling point.

32. Filtration: Separates solid substances from liquids and solutions.

33. Chromatography: Separates substances on the basis of differences in solubility in a solvent.

34. 1.4 Units of Measurement

35. SI Units Système International d’Unités Uses a different base unit for each quantity

36. Metric System Prefixes convert the base units into units that are appropriate for the item being measured.

37. Volume The most commonly used metric units for volume are the liter (L) and the milliliter (mL). □A liter is a cube 1 dm long on each side. □A milliliter is a cube 1 cm long on each side.

38. 1.5 Uncertainty in Measurements Different measuring devices have different uses and different degrees of accuracy.

39. Temperature: A measure of the average kinetic energy of the particles in a sample.

40. Temperature In scientific measurements, the Celsius and Kelvin scales are most often used. The Celsius scale is based on the properties of water. □0  C is the freezing point of water. □100  C is the boiling point of water.

41. Temperature The Kelvin is the SI unit of temperature. It is based on the properties of gases. There are no negative Kelvin temperatures. K =  C + 273.15

42. Temperature The Fahrenheit scale is not used in scientific measurements.  F = 9/5(  C) + 32  C = 5/9(  F − 32)

43. Density: Physical property of a substance d=d= mVmV

44. Uncertainty in Measurement

45. Significant Figures The term significant figures refers to digits that were measured. When rounding calculated numbers, we pay attention to significant figures so we do not overstate the accuracy of our answers.

46. Significant Figures 1.All nonzero digits are significant. 2.Zeroes between two significant figures are themselves significant. 3.Zeroes at the beginning of a number are never significant. 4.Zeroes at the end of a number are significant if a decimal point is written in the number.

47. Significant Figures When addition or subtraction is performed, answers are rounded to the least significant decimal place. When multiplication or division is performed, answers are rounded to the number of digits that corresponds to the least number of significant figures in any of the numbers used in the calculation.

48. Accuracy versus Precision Accuracy refers to the proximity of a measurement to the true value of a quantity. Precision refers to the proximity of several measurements to each other.

49. © 2012 Pearson Education, Inc. Dimensional Analysis We use dimensional analysis to convert one quantity to another. Most commonly, dimensional analysis utilizes conversion factors (e.g., 1 in. = 2.54 cm) 1 in. 2.54 cm 1 in. or

50. © 2012 Pearson Education, Inc. Dimensional Analysis Use the form of the conversion factor that puts the sought-for unit in the numerator: Given unit  desired unit desired unit given unit Conversion factor

51. © 2012 Pearson Education, Inc. Dimensional Analysis For example, to convert 8.00 m to inches, –convert m to cm –convert cm to in.  8.00 m  100 cm 1 m 1 in. 2.54 cm  315 in.

52. Ch. 2 Atoms, Molecules, and Ions

53. © 2012 Pearson Education, Inc. 2.1 Atomic Theory of Matter The theory that atoms are the fundamental building blocks of matter reemerged in the early nineteenth century, championed by John Dalton.

54. © 2012 Pearson Education, Inc. 2.1 Dalton's Postulates Each element is composed of extremely small particles called atoms.

55. © 2012 Pearson Education, Inc. 2.1 Dalton's Postulates All atoms of a given element are identical to one another in mass and other properties, but the atoms of one element are different from the atoms of all other elements.

56. © 2012 Pearson Education, Inc. 2.1 Dalton's Postulates Atoms of an element are not changed into atoms of a different element by chemical reactions; atoms are neither created nor destroyed in chemical reactions.

57. © 2012 Pearson Education, Inc. 2.1 Dalton's Postulates Compounds are formed when atoms of more than one element combine; a given compound always has the same relative number and kind of atoms.

58. © 2012 Pearson Education, Inc. 2.1 Law of Conservation of Mass The total mass of substances present at the end of a chemical process is the same as the mass of substances present before the process took place.

59. © 2012 Pearson Education, Inc. 2.2 The Electron Streams of negatively charged particles were found to emanate from cathode tubes, causing fluorescence. J. J. Thomson is credited with their discovery (1897).

60. © 2012 Pearson Education, Inc. 2.2 The Electron Thomson measured the charge/mass ratio of the electron to be 1.76  10 8 coulombs/gram (C/g).

61. © 2012 Pearson Education, Inc. 2.2 Millikan Oil-Drop Experiment Once the charge/mass ratio of the electron was known, determination of either the charge or the mass of an electron would yield the other.

62. © 2012 Pearson Education, Inc. 2.2 Millikan Oil-Drop Experiment Robert Millikan (University of Chicago) determined the charge on the electron in 1909.

63. © 2012 Pearson Education, Inc. 2.2 Radioactivity Radioactivity is the spontaneous emission of radiation by an atom. It was first observed by Henri Becquerel. Marie and Pierre Curie also studied it.

64. © 2012 Pearson Education, Inc. 2.2 Radioactivity Three types of radiation were discovered by Ernest Rutherford: –  particles –  particles –  rays

65. © 2012 Pearson Education, Inc. 2.2 The Atom, circa 1900 The prevailing theory was that of the “plum pudding” model, put forward by Thomson. It featured a positive sphere of matter with negative electrons imbedded in it.

66. © 2012 Pearson Education, Inc. 2.2 Discovery of the Nucleus Ernest Rutherford shot  particles at a thin sheet of gold foil and observed the pattern of scatter of the particles.

67. © 2012 Pearson Education, Inc. 2.2 The Nuclear Atom Since some particles were deflected at large angles, Thomson’s model could not be correct.

68. © 2012 Pearson Education, Inc. 2.2 The Nuclear Atom Rutherford postulated a very small, dense nucleus with the electrons around the outside of the atom. Most of the volume of the atom is empty space.

69. © 2012 Pearson Education, Inc. 2.3 Other Subatomic Particles Protons were discovered by Rutherford in 1919. Neutrons were discovered by James Chadwick in 1932.

70. © 2012 Pearson Education, Inc. 2.3 Subatomic Particles Protons and electrons are the only particles that have a charge. Protons and neutrons have essentially the same mass. The mass of an electron is so small we ignore it.

71. © 2012 Pearson Education, Inc. 2.3 Symbols of Elements Elements are symbolized by one or two letters.

72. © 2012 Pearson Education, Inc. 2.3 Symbols of Elements All atoms of the same element have the same number of protons, which is called the atomic number, Z.

73. © 2012 Pearson Education, Inc. 2.3 Symbols of Elements The mass of an atom in atomic mass units (amu) is the total number of protons and neutrons in the atom.

74. © 2012 Pearson Education, Inc. 2.3 Isotopes Isotopes are atoms of the same element with different masses. Isotopes have different numbers of neutrons.

75. © 2012 Pearson Education, Inc. 2.3 Atomic Mass Atomic and molecular masses can be measured with great accuracy using a mass spectrometer.

76. © 2012 Pearson Education, Inc. 2.3 Average Mass Because in the real world we use large amounts of atoms and molecules, we use average masses in calculations. Average mass is calculated from the isotopes of an element weighted by their relative abundances.

77. © 2012 Pearson Education, Inc. 2.4 Periodic Table The periodic table is a systematic catalog of the elements. Elements are arranged in order of atomic number.

78. © 2012 Pearson Education, Inc. 2.4 Periodic Table The rows on the periodic chart are periods. Columns are groups. Elements in the same group have similar chemical properties.

79. © 2012 Pearson Education, Inc. 2.4 Periodicity When one looks at the chemical properties of elements, one notices a repeating pattern of reactivities.

80. © 2012 Pearson Education, Inc. 2.4 Groups These five groups are known by their names.

81. © 2012 Pearson Education, Inc. 2.4 Periodic Table Nonmetals are on the right side of the periodic table (with the exception of H).

82. © 2012 Pearson Education, Inc. 2.4 Periodic Table Metalloids border the stair-step line (with the exception of Al, Po, and At).

83. © 2012 Pearson Education, Inc. 2.4 Periodic Table Metals are on the left side of the chart.

84. © 2012 Pearson Education, Inc. 2.5 Chemical Formulas The subscript to the right of the symbol of an element tells the number of atoms of that element in one molecule of the compound.

85. © 2012 Pearson Education, Inc. 2.5 Chemical Formulas Molecular compounds are composed of molecules and almost always contain only nonmetals.

86. © 2012 Pearson Education, Inc. 2.5 Diatomic Molecules These seven elements occur naturally as molecules containing two atoms: –Hydrogen –Nitrogen –Oxygen –Fluorine –Chlorine –Bromine –Iodine

87. © 2012 Pearson Education, Inc. 2.5 Types of Formulas Empirical formulas give the lowest whole- number ratio of atoms of each element in a compound. Molecular formulas give the exact number of atoms of each element in a compound.

88. © 2012 Pearson Education, Inc. 2.5 Types of Formulas Structural formulas show the order in which atoms are bonded. Perspective drawings also show the three-dimensional array of atoms in a compound.

89. © 2012 Pearson Education, Inc. 2.6 Ions When atoms lose or gain electrons, they become ions. –Cations are positive and are formed by elements on the left side of the periodic chart. –Anions are negative and are formed by elements on the right side of the periodic chart.

90. © 2012 Pearson Education, Inc. 2.6 Ionic Bonds Ionic compounds (such as NaCl) are generally formed between metals and nonmetals.

91. © 2012 Pearson Education, Inc. 2.6 Writing Formulas Because compounds are electrically neutral, one can determine the formula of a compound this way: –The charge on the cation becomes the subscript on the anion. –The charge on the anion becomes the subscript on the cation. –If these subscripts are not in the lowest whole- number ratio, divide them by the greatest common factor.

92. © 2012 Pearson Education, Inc. 2.6 Common Cations

93. © 2012 Pearson Education, Inc. 2.6 Common Anions

94. © 2012 Pearson Education, Inc. 2.7 Inorganic Nomenclature Write the name of the cation. If the anion is an element, change its ending to -ide; if the anion is a polyatomic ion, simply write the name of the polyatomic ion. If the cation can have more than one possible charge, write the charge as a Roman numeral in parentheses.

95. © 2012 Pearson Education, Inc. 2.7 Patterns in Oxyanion Nomenclature When there are two oxyanions involving the same element: –The one with fewer oxygens ends in -ite. –The one with more oxygens ends in -ate. NO 2 − : nitrite ; SO 3 2− : sulfite NO 3 − : nitrate; SO 4 2− : sulfate

96. © 2012 Pearson Education, Inc. 2.7 Patterns in Oxyanion Nomenclature Central atoms on the second row have bond to at most three oxygens; those on the third row take up to four. Charges increase as you go from right to left.

97. © 2012 Pearson Education, Inc. 2.7 Patterns in Oxyanion Nomenclature The one with the second fewest oxygens ends in -ite. –ClO 2 − : chlorite The one with the second most oxygens ends in -ate. –ClO 3 − : chlorate

98. © 2012 Pearson Education, Inc. Patterns in Oxyanion Nomenclature The one with the fewest oxygens has the prefix hypo- and ends in -ite. –ClO − : hypochlorite The one with the most oxygens has the prefix per- and ends in -ate. –ClO 4 − : perchlorate

99. © 2012 Pearson Education, Inc. Acid Nomenclature If the anion in the acid ends in -ide, change the ending to -ic acid and add the prefix hydro-. –HCl: hydrochloric acid –HBr: hydrobromic acid –HI: hydroiodic acid

100. © 2012 Pearson Education, Inc. 2.7 Acid Nomenclature If the anion in the acid ends in -ite, change the ending to -ous acid. –HClO: hypochlorous acid –HClO 2 : chlorous acid

101. © 2012 Pearson Education, Inc. 2.7 Acid Nomenclature If the anion in the acid ends in -ate, change the ending to -ic acid. –HClO 3 : chloric acid –HClO 4 : perchloric acid

102. © 2012 Pearson Education, Inc. 2.7 Nomenclature of Binary Compounds The less electronegative atom is usually listed first. A prefix is used to denote the number of atoms of each element in the compound (mono- is not used on the first element listed, however).

103. © 2012 Pearson Education, Inc. 2.7 Nomenclature of Binary Compounds The ending on the more electronegative element is changed to -ide. –CO 2 : carbon dioxide –CCl 4 : carbon tetrachloride

104. © 2012 Pearson Education, Inc. 2.7 Nomenclature of Binary Compounds If the prefix ends with a or o and the name of the element begins with a vowel, the two successive vowels are often elided into one. N 2 O 5 : dinitrogen pentoxide

105. © 2012 Pearson Education, Inc. 2.7 Nomenclature of Organic Compounds Organic chemistry is the study of carbon. Organic chemistry has its own system of nomenclature.

106. © 2012 Pearson Education, Inc. 2.7 Nomenclature of Organic Compounds The simplest hydrocarbons (compounds containing only carbon and hydrogen) are alkanes.

107. © 2012 Pearson Education, Inc. 2.7 Nomenclature of Organic Compounds The first part of the names just listed correspond to the number of carbons (meth- = 1, eth- = 2, prop- = 3, etc.).

108. © 2012 Pearson Education, Inc. 2.7 Nomenclature of Organic Compounds When a hydrogen in an alkane is replaced with something else (a functional group, like -OH in the compounds above), the name is derived from the name of the alkane. The ending denotes the type of compound. –An alcohol ends in -ol.

109. Which of the following is not a pure substance? a.water b.carbon dioxide c.carbon d. air

110. Solutions may be a.solids. b.liquids. c.gases. d. All of the above

111. Explain the difference between a scientific law (X) and a scientific theory (Y). a.X = is proven; Y = is not proven b.X = is not proven; Y = is proven c.X = tells what happens; Y = explains why things happen d.X = explains why things happen; Y = tells what happens

112. Properties that describe the way a substance reacts to form other substances are called a.physical properties. b.chemical properties. c.homogeneous properties. d.heterogeneous properties.

113. When nitric acid is mixed with copper metal, a brown gas forms. This is an example of a.an accident. b.a chemical reaction. c.a physical property. d.an extensive property.

114. A solution with a boiling point of 105 degrees Celsius contains either sugar or salt. How would you decide which is present? a.Distill the solution b.Filter the solution c.Use chromatography d. Taste the solution

115. Heat flows from an object a.at high temperature to an object at low temperature. b.at low temperature to an object at high temperature. c.to another object at the same temperature. d.at high elevation to an object at low elevation.

116. 120 milliliters = _____ fluid ounces. a.2 b.4 c.6 d.8

117. C = the speed of light = 3.00 x 10 8 meters/second = _____ furlongs/fortnight (8 furlongs = 1 mile) (14 days = 1 fortnight) a.7.5 x 10 10 b.1.3 x 10 11 c.2.2 x 10 11 d.1.8 x 10 12

118. In three trials, a student measures a solution’s concentration to be 0.010 M, 0.060 M, and 0.030 M. The accepted value is 0.034 M. The student’s data has a.good accuracy and good precision. b.good accuracy, but poor precision. c.poor accuracy, but good precision. d.poor accuracy and poor precision.

119. The measured quantity 0.082060 contains _____ significant figures. a.3 b.4 c.5 d.6

120. 6.03 g + 7.1 g = _____ g. a.13 b.13.1 c.13.13 d. 13.130

121. 7.1 g – 6.03 g = _____ g. a.1 b.1.1 c.1.07 d.1.070

122. 7.1 m x 6.03 m = _____ m 2. a.43 b.42.8 c.42.81 d.42.813

123. 6.03 g / 7.1 mL = _____ g/mL. a.0.8 b.0.85 c.0.849 d.0.849257

124. A temperature of –40 degrees Celsius is the same as (X) degrees Fahrenheit and (Y) Kelvins. a.X = –40, Y = 233 b.X = 233, Y = -40 c.X = 10, Y = 233 d.X = 233, Y = 10

125. The nucleus of an atom contains a.protons and neutrons. b.protons and electrons. c.electrons and neutrons. d.air.

126. Two atoms with the same atomic number but different mass numbers are called a.mutants. b.isomers. c.isotopes. d.symbiots.

127. The elements found on the left side of the periodic table tend to ______ electrons. a.gain b.lose c.keep d.share

128. A metal and a nonmetal react to form ________ compound. a.a molecular b.a mixed c.an empirical d.an ionic

129. Positive ions are called a.positrons. b.anions. c.cations. d.nucleons.

130. Compounds composed only of carbon and hydrogen are called a.binary acids. b.carbohydrates. c.hydrocarbons. d.alkanes.

131. The elements located in Group VIIA (Group 17) on the Periodic Table are called a.alkali metals. b.noble gases. c.chalcogens. d.halogens.

132. Acids produce _____ ions. a.OH -1 b.OH +1 c.H +1 d.H -1

133. Which of the formulas below does not represent a compound that actually exists? a.CaCO 3 b.H 2 O 2 c.KMnO 4 d.Na 2 PO 3

134. NaOCl is named a.sodium chlorate. b.sodium chlorite. c.sodium perchlorate. d.sodium hypochlorite.

135. LiNO 3 is named a.lithium nitrate. b.lanthanum nitrate. c.lanthanum nitrite. d.lithium nitrite.

136. Fe 2 O 3 is named a.diiron trioxide. b.iron(III) oxide. c.ferrous oxide. d.ironic oxide.

137. HIO 4 is named a.iodic acid. b.iodous acid. c.periodic acid. d.hydrogen iodate.

138. Cl 2 O 7 is named a.chlorine(VII) oxide. b.dichlorine hexaoxide. c.dichlorine heptaoxide. d.bichlorine heptaoxide.

139. C 3 H 8 is named a.ethane. b.propane. c.propanol. d.pentane.