STEM - GEN CHEM 1.pptxn.................

Lesson 1

1. GENERAL CHEMISTRY 1 SENIOR HIGH SCHOOL

2. 2 General Chemistry 1 Subject Description Composition, structure, and properties of matter; quantitative principles, kinetics, and energetics of transformations of matter; and fundamental concepts of organic chemistry S CIENCE T ECHNOLOGY E NGINEERING & M ATHEMATICS SPECIALIZED SUBJECT Grade Level : 11 Semester : 1 st /2 nd Hours/Semester : 80 Hrs.

3. 3 Contents UNIT 1 INTRODUCTION TO CHEMISTRY UNIT 2 CHEMICAL CALCULATION & REACTION UNIT 3 THE GASEOUS STATE OF MATTER UNIT 4 ATOMIC AND MOLECULARSTRUCTURES S CIENCE T ECHNOLOGY E NGINEERING & M ATHEMATICS SPECIALIZED SUBJECT STEM-ACADEMIC Grade Level : 11 Semester : 1 st /2 nd Hours/Semester : 80 Hrs.

4. UNIT 1 INTRODUCTION TO CHEMISTRY Lesson 1: Matter and Its Properties Lesson 2: Measurements Lesson 3: Atoms, Molecules, and Ions Lesson 4: Mole Concept 4

5. LESSON 1: MATTER AND ITS PROPERTIES OBJECTIVES OF THE DAY I will be able to describe the particulate nature of the different forms of matter; I will be able to classify the properties of matter; I will be able to differentiate pure substance and mixtures; elements and compounds; homogeneous and heterogeneous mixtures; 1 2 3 4 5 6 5

6. LESSON 1: MATTER AND ITS PROPERTIES OBJECTIVES OF THE DAY I will be able to recognize the formulas of some common substances; I will be able to discuss methods to separate the components of a mixtures; and I will be able to recognize chemical substances present in some consumer products 1 2 3 4 5 6 6

7. 7 MATTER Activity 1: What is Matter?

8. 8 Matter is anything that has mass and occupies space. Everything on earth has mass and takes up space .

9. 9 PARTICLES COMPOSING MATTER These are the smallest unit of matter that can’t be broken down chemically. These are groups of two or more atoms that are chemically bonded. These are particles that have gained or lost one or more of their valence electrons. ATOMS MOLECULES IONS

10. 10 STATES OF MATTER SOLID LIQUID GAS Activity 2: Table Completion

11. 11

12. 12

13. 13

14. 14

15. PLASMA THE 4 TH STATE OF MATTER It is a hot ionized gas consisting of approximately equal numbers of positively charged ions and negatively charged electrons. The characteristics of  plasmas  are significantly different from those of ordinary neutral gases so that  plasmas  are considered a distinct "fourth state of matter." 15

16. BOSE-EISTEIN CONDENSATE THE 5 TH STATE OF MATTER It is a state of matter in which separate atoms or subatomic particles, cooled to near absolute zero. When they reach that temperature the atoms are hardly moving relative to each other; they have almost no free energy to do so. At that point, the atoms begin to clump together, and enter the same energy states. 16

17. 17 PROPERTIES OF MATTER PHYSICAL PROPERTIES These can be measured and observed without changing the composition of the substance. CHEMICAL PROPERTIES These are the ability of a substance to react with other substances such as air, water, and base. INTENSIVE PROPERTIES It does not depend on the size or amount of the sample. EXTENSIVE PROPERTIES These can be affected by the size and amount of samples. According to changed involved during measurements of the property. According to dependence on amount of matter

18. 18

19. 19

20. 20

21. 21

22. 22

23. PHYSICAL PROPERTIES INTENSIVE PHYSICAL PROPERTIES EXTENSIVE PHYSICAL PROPERTIES Color Melting Point Density Solubility Conductivity Malleability Luster Viscosity Boiling Point Temperature Odor Mass Volume Length 23

24. CHEMICAL PROPERTIES 24 \ CHEMICAL PROPERTIES DESCRIPTION Combustibility Whether the substance undergoes combustion or not 2. Stability Whether the substance can be easily decomposed or not 3. Reactivity Whethe r it reacts with acids, bases, and oxygen, gas or not 4. Relative Activity Whether the material is more active or less active than other members of its chemical family 5. Ionization Whether it will break into charged particles when in solution with water or not. 6. Toxicity Whether substance can damage an organism or not.

25. Activity 3 Group the characteristics of the give substance according to their physical (extensive or intensive) or chemical properties. CHARACTERISTICS OF SOME SUBSTANCES PHYSICAL PROPERTIES CHEMICAL PROPERTIES INTENSIVE EXTENSIVE 1. The water in the container has a volume of 100 mL and a mass of 99.8 g. It is colorless , and tasteless. It has a density of 0.998g/mL, boils at 100 degrees Celsius, and freezes at 0 degree Celsius. It does not burn. It causes Iron to rust. 25

26. Activity 3 Group the characteristics of the give substance according to their physical (extensive or intensive) or chemical properties. CHARACTERISTICS OF SOME SUBSTANCES PHYSICAL PROPERTIES CHEMICAL PROPERTIES INTENSIVE EXTENSIVE 2. NaCl with a mass of 37.9 g is colorless , odorless , and salty solid crystals. It has melting point of 801 degree Celsius. When dissolved in 100 mL water, it conducts electricity. It reacts with silver nitrate to form a white precipitates. It also react with water to form chlorine gas, hydrogen gas, and sodium hydroxide. 26

27. 27 MATTER It is a matter that has a definite composition and distinct properties These are composed of two or more substances combined physically in various composition It is the simplest form of matter since it composed of only one kind of atom. It contains two or more kinds of atom chemically combined in definite proportion by mass  It is a solid, liquid, or gaseous  mixture  that has the same proportions of its components throughout any given sample. It is a mixture whose composition varies from one position to another within the sample. PURE SUBSTANCE MIXTURES ELEMENT COMPOUND HOMOGENEOUS MIXTURE HETEROGENOUS MIXTURE

28. Activity 4: Pure Substance or Mixture? 1. TABLE SUGAR 2. TABLE SALT 28

29. PURE SUBSTANCE OR MIXTURE? 3. IODIZED SALT 4. DISTILLED WATER 29

30. PURE SUBSTANCE OR MIXTURE? 4. SOFTDRINKS 5. OXYGEN GAS (TANK) 30

31. PURE SUBSTANCE OR MIXTURE? 6. BROWN SUGAR 7. HUMAN BREATH 31

32. Activity 5: HOMOGENEOUS OR HETEROGENEOUS? 1. RUBBING ALCOHOL 2. WATER &OIL 32

33. HOMOGENEOUS OR HETEROGENEOUS? 3. SALT & PEPPER 4. CARBONATED SOFTDRINKS 33

34. HOMOGENEOUS OR HETEROGENEOUS? 5. HUMAN BREATH 34

35. SEPERATING MIXTURES Chemist separate mixtures by using different methods. 35

36. SEPERATING MIXTURES Chemist separate mixtures by using different methods. Filtration is a process of separating the components of a suspension In Decantation the solid particles are allowed to settled first at the bottom and later, the liquid which is called supernatant is poured into another container leaving behind solid particle. Evaporation is the process of converting liquid to gas, is useful in sorting mixtures such as salt solution. Distillation is a process of separating a homogeneous mixture composed of two substances with different boiling points. 36

37. SEPERATING MIXTURES Chemist separate mixtures by using different methods. 5. Magnetic Separation is the process of separating elemental metals from other particles in a mixture. 6. Melting is a process that can be used in extricating mixture that contain two substances with different melting points. 7. Sublimation is a process of changing solid to gas without passing through the liquid state. 8. In Centrifugation, the mixture is poured into a special tube in the centrifuge apparatus, and is allowed to spin using centrifugal force. The spinning motion forces the sediments to settle at the bottom. The liquid can be poured off from the solid particles. 37

38. 9. Chromatography is another method of separating complex mixtures. It has various methods that can be used in separating mixture such as paper chromatography, which makes used of an adsorbent (filter paper or chromatogram paper), then separation depends upon the solubility of each component in the solvent. 38

39. PURE SUBSTANCES & MIXTURES IN CONSUMER PRODUCTS CONSUMER PRODUCT It is any item often bought for consumption. Convenience Product – those that appeal to a large segment of the market or those that are routinely bought. Household Cleaning Personal Care Product 39

40. PURE SUBSTANCES & MIXTURES IN CONSUMER PRODUCTS HOUSEHOLD CLEANING MATERIAL The most commonly used cleaning products are bleach, soaps, and detergents. These products have different compositions, specific uses, precautions for use, and costs. 40

41. PURE SUBSTANCES & MIXTURES IN CONSUMER PRODUCTS HOUSEHOLD CLEANING MATERIAL Bleach helps clean and whiten surfaces by generally lowering the stability of the chemical bonds in stain molecules. It can convert dirt into particles that can be easily washed away in conjunction with use of detergents. NaOCl (Sodium Hypochlorite) and H 2 O 2 (Hydrogen Peroxide) are most common bleaching agents that are strong oxidizers; they can burn then skin and eyes especially if used in concentrated form. 41

42. PURE SUBSTANCES & MIXTURES IN CONSUMER PRODUCTS HOUSEHOLD CLEANING MATERIAL Soap and Detergent are mixture of surfactants, water softeners, stain removers, enzymes and perfumes, among others. Surfactants render soaps and detergents capable of lowering the surface tension of water, which allows them to wet the surface to be cleaned. They also loosen and disperse water-insoluble solids making them washable with water. Soap and Detergents are generally not toxic and severely dangerous, but may cause irritation to the skin and eyes. 42

43. PURE SUBSTANCES & MIXTURES IN CONSUMER PRODUCTS PERSONAL CARE PRODUCTS Personal Care product constitute a diverse group of materials that improve the overall appearance of a person. These products are used to generally cleanse and beautify. Examples of highly demanded personal care products are makeup, lotions, and toothpastes. 43

44. LESSON 2: MEASUREMENTS OBJECTIVES OF THE DAY I will be able to describe the need for measurement; I will be able to carry out simple measurements of length, volume, and mass; and I will be able to differentiate the accuracy and the precision of a measurement 1 2 3 4 5 6 44

45. Keywords for the concepts to be learned; Measurements Unit of Measurements Accuracy Precision Significant figures Errors 45

46. Measurements The study of matter requires a certain degree of measurements , a process of determining the extent of the dimensions, quantity, or extent of something . Questions such as “How much….?” ,“How long…?” and “How many…?” simple cannot be answered without resorting to measurement. Q1. Can you cite some situations in daily life where a measurement is important? 46

47. Units of Measurements The most convenient system of units is the International System of Units (SI) . This system is the modern versions of metric system. 47

48. Units of Measurements The name of the fractional parts and the multiples of the base units are constructed by adding prefixes. These prefixes, shown in table, indicate the size of the unit relative to the base unit. 48

49. Uncertainty in Measurements A measured quantity contains some digits that are exactly known and one digit that is estimated. The estimated digit produces uncertainty in measurements. 49

50. Random Error and Systematic Error RANDOM ERROR (indeterminate error) is the uncertainty that arises from a scale reading which results from the uncontrolled variables in the measurement. It causes one measurement to differ slightly from the next. It comes from unpredictable changes during an experiment. Examples When weighing yourself on a scale, you position yourself slightly different each time. Measuring your height is affected by minor posture changes. 50

51. Random Error and Systematic Error SYSTEMATIC ERROR (determinate error) is the uncertainty that may come from a flaw in the equipment used or design of an experiment. These error are usually caused by measuring instruments that are incorrect calibrated or are used incorrect. Examples A worn out instrument An incorrectly calibrated or tared instrument A person consistently take an incorrect measurements 51

52. 52 Activity 6 Each dot is the result of a measurement whose value is indicated in the horizontal (or x-) axis. The plot presents the results of six measurements of the weight of a pebble whose true weight is 8.0 g. Determine whether each measurement is accurate or inaccurate, and precise or imprecise.

53. Precision and Accuracy Precision is the consistency of a result. If you measure a quantity several times and the values agrees closely with one another, then your measurement is precise.; however, if the values varied widely, then it is imprecise . Accuracy is determined when a certain quantitative value is relatively close to the “true value” 53

54. 54

55. Activity 7 55

56. Scientific Notation It is a simple way to write or keep track of very large or very small numbers without having to deal with a lot of zeros. It provides a convenient way of recording results and doing calculations. 56

57. 57

58. Significant Figures Significant figures are the digits in any measurement that are known certainty with an additional one digit which is uncertain. 58 RULES MEASURED NUMBERS NUMBER OF SIGNIFICANT All non zero digits are significant. 247 3 2. Zeroes between nonzero digits are significant. 20303 5 3. Zeroes to the left of the first nonzero digits are NOT significant 0.0200 3

59. Significant Figures 59 RULES MEASURED NUMBERS NUMBER OF SIGNIFICANT 4. If the number is less than 1, then only the zeros at the end of the number and the zero between nonzero digits are significant. 0.003560 4 5. If the number is greater than 1, then all the zeros written to the right of the decimal point are significant. 35600.00 7

60. Activity 9 Give the number of significant figures for each of the following measurements. 2 365 mm 309 cm 5.030 g/mL 0.0670 g 3.60 x 10 -4 60

61. Activity 10 Give the number of significant figures for each of the following measurements. 0.476 kg 89.7808 ft 0.430 mg 60.0 min 1 x 10 7 61

62. Rules for Significant Figures in Fundamental Operations In addition and subtraction, the answer must have the same number of decimal places as the measured number with the least number of decimal places. In multiplication and division, the answer must have the same number of significant figures as the measured number with the lowest number of significant figures. 62

63. Activity 11 Perform the following operations and write the answers in the proper number of significant figures. 4.87 m + 36.578 m + 4.34 m 8.9 mL ÷ 45 mL 68.980 cm – 67.16 cm 45.00 ft. x 3.00 ft. 14.4 g + 6.0 g 63

64. Rules in Rounding Off Oftentimes, the answers to computations contain too many insignificant digits. Hence it becomes necessary to round off numbers to attain the insignificant figures. Rounding off, therefore, is the process of removing, insignificant digits from calculated number. 64

65. Rules in Rounding Off The following rules should be applied to round off values to the correct number of digits. For a series of calculations, carry extra digits through to the final result, then round off. If the first digit to be deleted is…. a. 5 or greater, the last retained figure is increased by one b. 4 or less, the last retained figure is retained. 65

66. Conversion of Units (Dimensional Analysis) Dimensional Analysis is a process in which a conversion factor written in a form of ratio is used to change units given in the data to the units desired. The following are steps to be followed in doing dimensional analysis. Write the unknown quantity that is sought, including the units. Write all known conversion factors needed Begin with what is known and then multiply it by the identified conversion factor, cancelling similar units to get the unknown units. 66

67. METRIC AND ENGLISH CONVERSIONS QUANTITY METRIC ENGLISH CONVERSION MASS g, Kg lb , oz 1lb = 454 g 1kg = 2.2 lb 1 oz = 28.35g LENGTH cm, m, km in, ft , mi, Å 1 in = 2.54cm 1 m = 39.37 in 1 ft = 12 in 1 mi = 1.609 km 1 km = 0.62137 mi 1 Å = 10 -10 m VOLUME mL, L qt , pints, cups, tsp , tbsp , fl oz , gal 1 qt =946 mL 1 L = 1.057 qt 1 L = 2.12 pints 1 L = 4.23 cups 1 tsp = 4.93 mL 1tbsp = 14.79 mL 1 fl oz = 29.06 mL 1 gal = 3.79 L 67

68. Activity 12 Sample Problems! The lemon juice drink contains 500.0 mg of vitamin C. Express the vitamin C content in grams. 2. A fitness drink measures 0.300 L. Express the volume in L 3. Calculate the number of centimeters in 53.5 inches. 68

69. Density Measurement Density measurement is one of the common measurements done in the laboratory. It involves getting Mass, Volume and Temperature of an Object. Mass is the quantity of matter in the object. It is determined by weighing the object, using balance. The SI basic unit of mass is the Kilogram, but the gram is more convenient to use. 69

70. Density Measurement B. Volume is the amount of space occupied by a substance. In liquids, the volume can be determined using a graduated cylinder, while solids, the volume can be determined by two methods. For regularly shaped-solids , the volume formula for the particular shaped is used. Some formula that may be used are the following: Rectangular solid = L x W x H Cylindrical Solid = π r 2 h Cubic solid = S x S x S Spherical solid = 4/3 π r 3 70

71. Density Measurement 2. For irregularly-shaped solids, the water displacement method is used. C. Temperature tells how hot or cold an object is. It is commonly marked either by o C (Celsius) or o F (Fahrenheit); although the SI basic unit for temperature is the K (Kelvin). To convert o C or o F to K, the following are used. K = o C + 273.15 K = ( o F + 459.67) x 5/9 71

72. Density Measurement DENSITY is the ratio of the mass of an object to the volume it occupies. DENSITY =   72

73. Activity 13 Sample Problems A sample amount of sugar has a mass of 250.0 g and a volume of 157.3 cm 3. What is its density in grams per cubic centimeter ? 2. Gold metal has a density of 19.3 g/cm 3 . What is the volume in cubic centimeter of a 500.0 g bar of gold metal? 73

74. Activity 14 Sample Problem 1. The volume and the mass of two objects (A & B) have been obtained in order to determine their densities, respectively. Identify which object is denser. 74 OBJECT METHOD USED FOR DETERMINING THE VOLUME MASS A By measuring its dimension L = 2.0 cm, W = 2.5 cm H = 15 cm 90.0 g B By water displacement method: Final Volume (Water + object) = 100 mL Initial Volume (Water) = 80.0 mL 65.0 g

75. Lesson 3: Atoms, Molecules and Ions OBJECTIVES OF THE DAY! I will be able to describe and discuss the basic laws of chemical change; I will be able to discuss how Dalton’s Atomic Theory could explain the basic laws of chemical changes; I will be able to give the information provided by the atomic number and mass number of an atom and its isotopes 75

76. Lesson 3: Atoms, Molecules and Ions OBJECTIVES OF THE DAY! I will be able to differentiate atoms, molecules, and ions; I will be able to write the chemical formula of some molecules; I will be able to differentiate a molecular formula and an empirical formula; and I will be able to give the name of a compound, given its chemical formula. 76

77. Keywords a. Law of Conservation of Matter h. Law of Definite Proportion b. Law of Multiple Proportion i . Dalton’s Atomic Theory c. Atomic number j. Mass number d. Isotope k. Atom e. Molecule l. Ion f. Chemical formula m. Molecular formula g. Empirical formula 77

78. LAWS OF CHEMICAL CHANGE These laws were inferred from several experiments conducted during the 18th century using a balance for the measurements: Law of Conservation of Mass Law of Definite Proportion Law of Multiple Proportion 78

79. A. Law of Conservation of Mass ANTOINE LAVOISIER , a brilliant French chemist, formulated this law by describing one of his experiments involving mercuric oxide. He placed a small amount of mercuric oxide, a red solid, inside a retort and sealed the vessel tightly. 79

80. A. Law of Conservation of Mass He weighed the system, and then subjected it to high temperature. During the heating, the red solid turned into a silvery liquid. This observation indicated that a chemical reaction took place. After which, the setup was cooled and then weighed. The weight of the system was found to be the same as before heating. 80

81. A. Law of Conservation of Mass In a chemical reaction, no change in mass takes place. The total mass of the products is equal to the total mass of the reactant. 81

82. B. Law of Definite Proportion A compound always contains the same constituent elements in a fixed or definite proportion by mass. If water samples coming from different sources are analyzed, all the samples will contain the same ratio by mass of hydrogen to oxygen. 82

83. Sample Problems A pure sample of Sodium Flouride ( NaF ) contains 35g of Sodium. How many grams of Flourine are present in this sample? If there are 42g of H in a sample of pure Methane (CH 4 ), How many grams of Carbon are present? If there are 19g of oxygen in a sample of Aluminum Oxide (Al 2 O 3 ), How many grams of Aluminum are present? 83

84. C. Law of Multiple Proportions If two elements can combine to form more than one compound, the masses of one element that will combine with a fixed mass of the other element are in a ratio of small whole numbers. 84

85. Dalton’s Atomic Theory In 1808, John Dalton published his book A New System of Chemical Philosophy, where he proposed an atomic theory of matter that can explain chemical observations as predicted by the three fundamental laws. 85

86. Dalton’s Atomic Theory The atomic theory comprised the following postulates: 1. Matter is made up of extremely small indivisible particles called atoms. 86

87. Dalton’s Atomic Theory The atomic theory comprised the following postulates: 2. Atoms of the same element are identical, and are different from those of other elements. 87

88. Dalton’s Atomic Theory The atomic theory comprised the following postulates: 3. Compounds are composed of atoms of more than one element, combined in definite ratios with whole number values. 88

89. Dalton’s Atomic Theory The atomic theory comprised the following postulates: 4. During a chemical reaction, atoms combine, separate, or rearrange. No atoms are created and no atoms disappear. 89

90. During the time of Dalton, the atom was believed to be the smallest particle comprising substances. However, before the end of the 19th century, experiments provided proof of the existence of smaller particles within the atom. 90

91. Activity 15 Recall the particles contained in an atom (or the subatomic particles ) and differentiate the particles in terms of location, charge, and relative mass by filling up the following table: 91

92. Activity 15 Recall the particles contained in an atom (or the subatomic particles ) and differentiate the particles in terms of location, charge, and relative mass by filling up the following table: 92

93. Atomic Number and Mass Number An atom of an element may be represented in a certain configuration that includes its atomic number (Z) and Mass number (A) , written as the left superscript and left subscript, respectively of the element symbol.   93 mass number (A) atomic number (Z) Symbol of Element

94. Atomic Number and Mass Number The atomic number of an element represents the number of protons in its nucleus. Because an atom as a whole is electrically neutral, the atomic number also specifies the number of electron present. ATOMIC NUMBER = NUMBER OF PROTONS = NUMBER OF ELECTRONS 94

95. Atomic Number and Mass Number The mass number of an atom is the sum of the number of protons and neutrons in its nucleus. Thus, the mass number gives the number of subatomic particles present in the nucleus. MASS NUMBER = NUMBER OF PROTONS + NUMBER OF NEUTRONS 95

96. Activity 16 COMPLETE THE TABLE BELOW 96

97. Isotopes are atoms of an element having the same atomic number but different mass number. The existence of isotopes was shown by mass spectroscopy experiments, wherein elements were found to be composed of several types of atoms, each with different masses. a. The atomic number identifies an element. The atoms of isotopes of an element have the same number of protons and electrons. b. The atoms of isotopes of an element differ in the number of neutrons. 97

98. Atoms, Molecules and Ions 98

99. Atoms, Ions and Molecules Of all the elements, only six exist as single atoms, namely Helium, Neon, Argon, Krypton, Xenon and Radon . Most matters are composed of ions formed from atoms. A molecule is a combination of at least two atoms in a definite proportion, bound together by covalent bonds. 99

100. Ions When a neutral atom gain or loses one or more electrons, it becomes an electrically charged particles called ion. 100

101. Ions Metals tend to lose electrons and become positively charged cations . Nonmetals , on the other hand, gain electrons and become negatively charged anions. The number of electron lost or gained is the charged number. Ions can be made up of only one atom (monoatomic) or more than one type of atom (polyatomic). 101

102. 102

103. 103

104. Naming Monoatomic Ions Monoatomic ions are named based on the element. For cations , the name of the element is unchanged. If an element can form two ions of different charges, the name, which is usually derived from its Latin name, is modified by the suffix – ic for the ion with the higher charge, and – ous for that with the lower charge. 104

105. 105

106. 106 Naming Monoatomic Ions b. The monoatomic anions are named by attaching the suffix –ide to the first few letters (root) of nonmetal name,

107. Activity 17.1 Name the following cations below. 107

108. Activity 17.2 Name the following anions below. 108

109. 109 Several anions are polyatomic and are named based on the atomic constituents and the suffix – ide . The most common examples are: a. OH- – hydroxide ion b. CN- – cyanide ion

110. 110 A number of polyatomic anions containing oxygen atoms are named based on the root word of the central (or non-oxygen) atom and the suffix –ate for the one with more oxygen atoms and – ite for the one with less oxygen atom. Some anions have common names ending with the suffix –ate .

111. 111

112. 112 Chemical Formula The composition of a molecule or an ion can be represented by a chemical formula. The formula consists of the symbols of the atoms making up the molecule. If there is more than one atom present, a numerical subscript is used.

113. 113 Chemical Formula There are two types of chemical formula - the molecular formula and empirical formula.

114. Chemical Formula Molecular Formula indicates the actual number of each element in a compound. Emperical Formula is the simplest chemical formula. It only shows relative ratio between the number of atoms of the different elements present in the compound. 114

115. Activity 18 Write the empirical formula of the following molecules. C 2 H 4 O 2 C 8 H 12 N 4 C 8 H 10 P 4 O 10 PH 3 115

116. Naming Compounds A. IONIC COMPOUNDS ( cation and anion) For Binary Compounds, metal cations take their names from the elements, while the anion take the first part of the name of element , and add the suffix – ide end. 116 CATION ANION COMPOUND NAME OF COMPOUND Na + O -2 Na 2 O Sodium oxide Mg +2 N -3 Mg 3 N 2 Magnesium nitride Al +3 O -2 Al 2 3 Aluminum Oxide

117. Naming Compounds 2. For Ternary Compounds, the cation goes first in its name before the polyatomic ion which usually ends with – ite or -ate 117 CATION ANION COMPOUND NAME OF COMPOUND Na + NO 3 -1 Na NO 3 Sodium Nitrate Na + NO 2 -1 Na NO 2 Sodium Nitrite

118. Naming Compounds 3. For compounds containing a metallic ion of variable charge, either the classical method or the stock method of naming may be used. In the classical method , the name of metallic ions ends in – ous (for lower charge) and – ic (for higher charge) In the stock method , the metal is named first followed by the value of the charge written in roman numeral (enclosed in parenthesis) 118

119. Naming Compounds B. MOLECULAR COMPOUNDS (TWO NONMETALS) For one pair of elements that form several different compounds, Greek prefixes are used to determine the number of each element in the compound. For the first element, the prefixes “mono” is omitted. Examples CO – carbon monoxide CO 2 – carbon dioxide N 2 O 4 – dinitrogen tetraoxide 119

120. 120

121. Naming Compounds B. MOLECULAR COMPOUNDS (TWO NONMETALS) For binary compounds, place the name of the first element; then follow it with the second element. The second element is named by adding –ide to the root of the element name. Examples: HCl – hydrogen chloride HBr – hydrogen bromide 121

122. Naming Compounds B. MOLECULAR COMPOUNDS (TWO NONMETALS) For binary compounds considered as acids, use the prefix hydro- followed it with second element. The second element is named by adding –ide to the root of the element name. Examples a. HCl – hydrogenchloric acid b. HBr – hydrobromic acid 122

123. Naming Compounds B. MOLECULAR COMPOUNDS (TWO NONMETALS) Oxy-acids, those that contain hydrogen, oxygen and another element, is named in two ways – For anions ending with – ate, change –ate to – ic ; then, follow it with the word acid. For anions ending with – ite , change – ite to – ous ; then follow it with the word acid. 123

124. Lesson 4: Mole Concept OOTD: OBJECTIVES OF THE DAY At the end of this lesson, the students must be able to explain the meaning of average atomic mass define a mole; illustrate Avogadro’s number with examples; determine the molar mass of elements and compounds; calculate the mass of a given number of moles of an element or compound, or vice versa; and calculate the mass of a given number of particles of an element or compound, or vice versa. 124

125. PERFORMANCE TASK THIRD QUARTER – GENERAL CHEMISTRY 1 History of the Atomic Structure J.J Thomson’s Plum Pudding Model John Dalton’s Billiard Ball Model E. Rutherford’s Nuclear Model Niels Bohr’s Planetary Model Schrodinger & Heisenberg Quantum Mechanical Model 125

126. Ms. Lilia sells shelled peanuts in a store. But she meets customers asking for 150 peanuts, another for 750 peanuts, and another for 2,000 peanuts. Obviously, it will take Ms. Lilia a very long time to count the peanuts. What would be another way to count them ? 126

127. 127

128. Atomic Mass and Atomic Mass Unit Question: Is it possible to use the same procedure to count atoms. Why or why not? Whether it is peanuts or mongo beans or candies or atoms, the procedure should be the same. The problem, however, is atoms are very, very small and it is not possible to see them and count them individually to get the average mass. We need to look for another way to get the average mass of the atom. 128

129. Atomic Mass and Atomic Mass Unit Experiments have shown that atoms have different masses relative to one another. For example, a Mg atom is experimentally reported to be twice as heavy as a carbon atom ; a silicon atom is twice the mass of a nitrogen atom. It is possible to make a relative scale if one atom is chosen as the reference or standard atom against which the masses of the other atoms are measured. 129

130. Atomic Mass and Atomic Mass Unit By international agreement, the reference atom chosen is the C-12 isotope which contains six protons and six neutrons. By definition, one atom of C-12 has a mass of exactly 12 atomic mass units ( amu ). One amu , therefore, is one-twelfth (1/12) the mass of a C-12 atom. 130

131. Atomic Mass and Atomic Mass Unit Example The atomic mass of Cu-63 is 62.93 amu . This means that relative to C-12, one atom of Cu-63 is 62.93/12 or 5.244 times the mass of a C-12 atom. Try This! One atom of Se-77 is 6.410 times as heavy as an atom of C-12. What is the atomic mass of Se-77? 131

132. Average Atomic Mass The atomic mass of the atoms of an element is the average atomic masses of the naturally occurring isotopes of this element. The periodic table provides the average atomic mass which takes into account the different isotopes of an element and their relative abundances. NOTE: It is not a simple average that is taken but a weighted average 132 For example The average atomic mass of Oxygen is 15.999, not 16.00 The 15.999 is calculated by considering the naturally-occurring isotopes of Oxygen, namely Oxygen-16, Oxygeny-17 and Oxygen-18

133. Average Atomic Mass Average atomic masses are obtained by multiplying the mass of an isotopes by its fractional abundance, as shown as follows. Isotopes of elements occur in different abundances. Some are more abundant than others. Chlorine has two isotopes. The natural abundance of Cl-35 is 75% while that of Cl-37 is 25%. This means that if you have 100 atoms of chlorine, 75 of them will be Cl-35 and 25 of them will be Cl-37. 133 ELEMENT MASS NUMBER ISOTOPIC MASS % ABUNDANCE AVERAGE ATOMIC MASS Oxygen 16 15.9949 u 99.76% 15.999 u 17 16.9991 u 0.04% 18 17.9992 u 0.20%

134. 2. Magnesium , on the other hand, has three isotopes with varying abundances: Mg-24,Mg-25, and Mg-26, 11.01 have 78.99%, 10.00%, and 11.01% abundance, respectively. 3. For carbon , the natural abundance of C-12 is 98.90% while that of C-13 is 1.10%. The atomic mass of C-13 has been determined to be 13.00335 amu while that of C-12 is exactly 12 amu . Relative Atomic Mass is the ratio of the average atomic mass of an atom to one atomic mass unit ( amu ) Hence, its value is similar with average atomic mass, except that it has no unit. 134

135. Activity 19 Copper has two stable isotopes with the following masses and % abundances: Cu-63 (62.93 amu , 69.09% abundance) and Cu-65 (64.9278 amu , 30.91% abundance). Calculate the average atomic mass of copper. 135

136. The Avogadro’s Number In the SI system, the mole (mole) is defined as the amount of substance containing the same number of particles as there are atoms in exactly 12 g of carbon-12 isotope. One mole of a substance is equivalent to the Avogdro’s number of particles 6.02 x 10 23 This number is so-named in honor of the Italian scientist, Amadeo Avogadro 136

137. The Avogadro’s Number Thus, based from the definition. It can be said that One mole of an element is numerically equal to its atomic mass unit. One mole of an element contains 6.02 x 10 23 atoms One mole of molecular compound contains 6.02 x 10 23 molecules One mole of ionic compound contains 6.02 x 10 23 cations and 6.02 x 10 23 anions 137

138. Molar Mass The molar mass of a compound (molecular or ionic) is the mass in grams of one mole of a substance. It is numerically equal to the sum of the masses of the elements (in amu ) that make up the compound. The molar mass is obtained by multiplying the number of atoms by the atomic mass of each element, and getting the sum. The unit for molar mass is g/mol. 138

139. Molar Mass The molar mass of a compound (molecular or ionic) is the mass in grams of one mole of a substance. It is numerically equal to the sum of the masses of the elements (in amu ) that make up the compound. The molar mass is obtained by multiplying the number of atoms by the atomic mass of each element, and getting the sum. The unit for molar mass is g/mol. 139

140. Activity 20 Calculate the molar mass of the following compounds C 3 H 5 N 3 O 9 (NH 2 ) 2 CO Hg(OCN) 2 140

141. Formula Mass and Molecular Mass Formula Mass is used for compounds that exists as ions, such as NaCl . It is expressed in amu or u, and is numerically equal to the molar mass expressed in grams per mole of a substance. Molecular Mass is used for compounds that exist as molecules, such as water (H 2 O) It is numerically equal to the molar mass and has a unit amu . 141 IONIC COMPOUND COMMON NAME MOLAR MASS FORMULA MASS NaCl Table Salt 58g/ mol 58 amu CaO Quicklime 56g/mole 56 amu MOLECULE COMMON NAME MOLAR MASS MOLECULAR MASS CO 2 Dry Ice 44g/ mol 44 amu C 12 H 22 O 11 Dextrose 342g/mole 342 amu

142. Calculation Involving Formulas The Avogadro’s number and molar mass are important to enable conversions between mass and moles of atoms or molecules, ion and vice versa. The following are the conversion factors that can be used in calculations involving formulas. Where X represents the symbol of atoms, ions, or the formula of the compound. and   142

143. Sample Problems A. Conversion between atoms, molecules, or ions and mass Zinc is an essential mineral that is naturally occurring found in foods and is also available as dietary supplement. How many atoms are in 16.5 g of Zinc? 2. Ammonia (NH 3 ) is used for fertilizers and many other things. How many molecules of ammonia are present in 0.334 g of ammonia? 143

144. Sample Problems B. Conversion between mass and moles Ammonium Nitrate (NH 4 NO 3 ) is a main component of explosive mixtures used in mining, quarrying, and civil construction. If an explosive contains 345.0 g of ammonium nitrate, how many mole of ammonium nitrate are present in the explosive? Copper is used for the absorption and used of iron in the formation of haemoglobin. How many grams of Cu are present in 3.87 mol copper? 144