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Chemical Compounds Section 1 Measuring Matter Section 2 Mass and the Mole Section 3 Moles of Compounds Section 4 Empirical and Molecular Formulas Section 5 Oxidation States Section 6 Naming Compounds Exit Section1 Measuring Matter • Explain how a mole is used to indirectly count the number of particles of matter. molecule: two or more atoms that covalently bond together to form a unit • Relate the mole to a common everyday counting unit. mole • Convert between moles and number of representative particles. Avogadro’s number Chemists use the mole to count atoms, molecules, ions, and formula units. Counting Particles • Chemists need a convenient method for accurately counting the number of atoms, molecules, or formula units of a substance. • The mole is the SI base unit used to measure the amount of a substance. • 1 mole is the amount of atoms in 12 g of pure carbon-12, or 6.02 1023 atoms. • The number is called Avogadro’s number. Converting Between Moles and Particles • Conversion factors must be used. • Moles to particles Number of molecules in 3.50 mol of sucrose Converting Between Moles and Particles (cont.) • Particles to moles • Use the inverse of Avogadro’s number as the conversion factor. Section1 Assessment What does the mole measure? A. mass of a substance B. amount of a substance C. volume of a gas D C A 0% B D. density of a gas A. A B. B C. C 0% 0% 0% D. D Section1 Assessment What is the conversion factor for determining the number of moles of a substance from a known number of particles? A D. 1 mol 6.02 1023 particles 0% D C. 1 particle 6.02 1023 C B. A. A B. B C. C 0% 0% 0% D. D B A. Section 2 Mass and the Mole • Relate the mass of an atom conversion factor: a to the mass of a mole of ratio of equivalent atoms. values used to express the same quantity in • Convert between number different units of moles and the mass of an element. • Convert between number of moles and number of atoms of an element. molar mass A mole always contains the same number of particles; however, moles of different substances have different masses. The Mass of a Mole • 1 mol of copper and 1 mol of carbon have different masses. • One copper atom has a different mass than 1 carbon atom. The Mass of a Mole (cont.) • Molar mass is the mass in grams of one mole of any pure substance. • The molar mass of any element is numerically equivalent to its atomic mass and has the units g/mol. Using Molar Mass • Moles to mass 3.00 moles of copper has a mass of 191 g. Using Molar Mass (cont.) • Convert mass to moles with the inverse molar mass conversion factor. • Convert moles to atoms with Avogadro’s number as the conversion factor. Using Molar Mass (cont.) • This figure shows the steps to complete conversions between mass and atoms. Section2 Assessment The mass in grams of 1 mol of any pure substance is: A. molar mass B. Avogadro’s number D A 0% C D. 1 g/mol A. A B. B C. C 0% 0% 0% D. D B C. atomic mass Section 2 Assessment Molar mass is used to convert what? A. mass to moles B. moles to mass C. atomic weight D C A 0% B D. particles A. A B. B C. C 0% 0% 0% D. D Section 3 Moles of Compounds • Recognize the mole relationships shown by a chemical formula. • Calculate the molar mass of a compound. • Convert between the number of moles and mass of a compound. • Apply conversion factors to determine the number of atoms or ions in a known mass of a compound. representative particle: an atom, molecule, formula unit, or ion Chemical Formulas and the Mole • Chemical formulas indicate the numbers and types of atoms contained in one unit of the compound. • One mole of CCl2F2 contains one mole of C atoms, two moles of Cl atoms, and two moles of F atoms. The Molar Mass of Compounds • The molar mass of a compound equals the molar mass of each element, multiplied by the moles of that element in the chemical formula, added together. • The molar mass of a compound demonstrates the law of conservation of mass. Converting Moles of a Compound to Mass • For elements, the conversion factor is the molar mass of the compound. • The procedure is the same for compounds, except that you must first calculate the molar mass of the compound. Converting the Mass of a Compound to Moles • The conversion factor is the inverse of the molar mass of the compound. Converting the Mass of a Compound to Number of Particles • Convert mass to moles of compound with the inverse of molar mass. • Convert moles to particles with Avogadro’s number. Converting the Mass of a Compound to Number of Particles (cont.) • This figure summarizes the conversions between mass, moles, and particles. Section 3 Assessment How many moles of OH— ions are in 2.50 moles of Ca(OH)2? A. 2.00 B. 2.50 D A 0% C D. 5.00 A. A B. B C. C 0% 0% 0% D. D B C. 4.00 Section 3 Assessment How many particles of Mg are in 10 moles of MgBr2? A. 6.02 1023 B. 6.02 1024 D A 0% C D. 1.20 1025 A. A B. B C. C 0% 0% 0% D. D B C. 1.20 1024 Section 4 Empirical and Molecular Formulas • Explain what is meant by the percent composition of a compound. • Determine the empirical and molecular formulas for a compound from mass percent and actual mass data. percent by mass: the ratio of the mass of each element to the total mass of the compound expressed as a percent percent composition empirical formula molecular formula A molecular formula of a compound is a whole-number multiple of its empirical formula. Percent Composition • The percent by mass of any element in a compound can be found by dividing the mass of the element by the mass of the compound and multiplying by 100. Percent Composition (cont.) • The percent by mass of each element in a compound is the percent composition of a compound. • Percent composition of a compound can also be determined from its chemical formula. Empirical Formula • The empirical formula for a compound is the smallest whole-number mole ratio of the elements. • You can calculate the empirical formula from percent by mass by assuming you have 100.00 g of the compound. Then, convert the mass of each element to moles. • The empirical formula may or may not be the same as the molecular formula. Molecular formula of hydrogen peroxide = H2O2 Empirical formula of hydrogen peroxide = HO Molecular Formula • The molecular formula specifies the actual number of atoms of each element in one molecule or formula unit of the substance. • Molecular formula is always a whole-number multiple of the empirical formula. Molecular Formula (cont.) Chemical Composition Halothane C2HBrClF3 Mole ratio nC/nhalothane Mass ratio mC/mhalothane M(C2HBrClF3) = 2MC + MH + MBr + MCl + 3MF = (2 12.01) + 1.01 + 79.90 + 35.45 + (3 19.00) = 197.38 g/mol Example Calculating the Mass Percent Composition of a Compound Calculate the molecular mass M(C2HBrClF3) = 197.38 g/mol For one mole of compound, formulate the mass ratio and convert to percent: (2 12.01) g %C 100% 12.17% 197.38 g Example (2 12.01) g %C 100% 12.17% 197.38 g 1.01g %H 100% 0.51% 197.38 g 79.90 g % Br 100% 40.48% 197.38 g 35.45 g %Cl 100% 17.96% 197.38 g (3 19.00) g %F 100% 28.88% 197.38 g Empirical formula 5 Step approach: 1. 2. 3. 4. 5. Choose an arbitrary sample size (100g). Convert masses to amounts in moles. Write a formula. Convert formula to small whole numbers. Multiply all subscripts by a small whole number to make the subscripts integral. Example Determining the Empirical and Molecular Formulas of a Compound from Its Mass Percent Composition. Dibutyl succinate is an insect repellent used against household ants and roaches. Its composition is 62.58% C, 9.63% H and 27.79% O. Its experimentally determined molecular mass is 230 u. What are the empirical and molecular formulas of dibutyl succinate? Step 1: Determine the mass of each element in a 100g sample. C 62.58 g H 9.63 g O 27.79 g Example Step 2: Convert masses to amounts in moles. 1 mol C 5.210 mol C 12.011 g C 1 mol H nH 9.63 g H 9.55 mol H 1.008 g H 1 mol O nO 27.79 g O 1.737 mol O 15.999 g O nC 62.58 g C Step 3: Write a tentative formula. C5.21H9.55O1.74 Step 4: Convert to small whole numbers. C2.99H5.49O Example Step 5: Convert to a small whole number ratio. Multiply 2 to get C5.98H10.98O2 The empirical formula is C6H11O2 Step 6: Determine the molecular formula. Empirical formula mass is 115 u. Molecular formula mass is 230 u. The molecular formula is C12H22O4 Oxidation States Metals tend to lose electrons. Non-metals tend to gain electrons. Na Na+ + e- Cl + e- Cl- Reducing agents Oxidizing agents We use the Oxidation State to keep track of the number of electrons that have been gained or lost by an element. Rules for Oxidation States 1. The oxidation state (OS) of an individual atom in a free element is 0. 2. The total of the OS in all atoms in: i. Neutral species is 0. ii. Ionic species is equal to the charge on the ion. 3. In their compounds, the alkali metals and the alkaline earths have OS of +1 and +2 respectively. 4. In compounds the OS of fluorine is always –1 Rules for Oxidation States 5. In compounds, the OS of hydrogen is usually +1 6. In compounds, the OS of oxygen is usually –2. 7. In binary (two-element) compounds with metals: i. Halogens have OS of –1, ii. Group 16 have OS of –2 and iii. Group 15 have OS of –3. Example Assigning Oxidation States. What is the oxidation state of the underlined element in each of the following? a) P4; b) Al2O3; c) MnO4-; d) NaH a) P4 is an element. P OS = 0 b) Al2O3: O is –2. O3 is –6. Since (+6)/2=(+3), Al OS = +3. c) MnO4-: net OS = -1, O4 is –8. Mn OS = +7. d) NaH: net OS = 0, rule 3 beats rule 5, Na OS = +1 and H OS = -1. Naming Compounds Trivial names are used for common compounds. A systematic method of naming compounds is known as a system of nomenclature. Inorganic compounds Inorganic Nomenclature Binary Compounds of Metals and Nonmetals NaCl = electrically neutral sodium chloride name is unchanged “ide” ending MgI2 = magnesium iodide Al2O3 = aluminum oxide Na2S = sodium sulfide Binary Compounds of Two Non-metals Molecular compounds usually write the positive OS element first. HCl hydrogen chloride Some pairs form more than one compound mono 1 penta 5 di 2 hexa 6 tri 3 hepta 7 tetra 4 octa 8 Binary Acids Acids produce H+ when dissolved in water. They are compounds that ionize in water. Emphasize the fact that a molecule is an acid by altering the name. HCl hydrogen chloride hydrochloric acid HF hydrogen fluoride hydrofluoric acid Polyatomic Ions Polyatomic ions are very common. Table 3.3 gives a list of some of them. Here are a few: ammonium ion NH4+ acetate ion C2H3O2- carbonate ion CO32- hydrogen carbonate HCO3- hypochlorite ClO- phosphate PO43- chlorite ClO2- hydrogen phosphate HPO42- chlorate ClO3- sulfate SO42- perchlorate ClO4- hydrogensulfate HSO4- •Slide 49 of 37 •General Chemistry: Chapter 3 •Prentice-Hall © 2002 How many moles of hydrogen atoms are in one mole of H2O2? A. 1 B. 2 D A 0% C D. 0.5 A. A B. B C. C 0% 0% 0% D. D B C. 3 How many moles of Al are in 2.0 mol of Al2Br3? A. 2 B. 4 D A 0% C D. 1 A. A B. B C. C 0% 0% 0% D. D B C. 6 How many atoms of hydrogen are in 3.5 mol of H2S? A. 7.0 1023 B. 2.1 1023 D A 0% C D. 4.2 1024 A. A B. B C. C 0% 0% 0% D. D B C. 6.0 1023 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. Atoms, Molecules, and Ions •© 2009, Prentice-Hall, Patterns in Oxyanion Nomenclature • When there are two oxyanions involving the same element: – The one with fewer oxygens ends in -ite. • NO2− : nitrite; SO32− : sulfite – The one with more oxygens ends in -ate. • NO3− : nitrate; SO42− : sulfate Atoms, Molecules, and Ions •© 2009, Prentice-Hall, Patterns in Oxyanion Nomenclature • The one with the second fewest oxygens ends in -ite. – ClO2− : chlorite • The one with the second most oxygens ends in -ate. – ClO3− : chlorate Atoms, Molecules, and Ions •© 2009, Prentice-Hall, Patterns in Oxyanion Nomenclature • The one with the fewest oxygens has the prefix hypoand ends in -ite. – ClO− : hypochlorite • The one with the most oxygens has the prefix per- and ends in -ate. – ClO4− : perchlorate Atoms, Molecules, and Ions •© 2009, Prentice-Hall, 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 Atoms, Molecules, and Ions •© 2009, Prentice-Hall, Acid Nomenclature • If the anion in the acid ends in -ite, change the ending to -ous acid. – HClO: hypochlorous acid – HClO2: chlorous acid Atoms, Molecules, and Ions •© 2009, Prentice-Hall, Acid Nomenclature • If the anion in the acid ends in -ate, change the ending to -ic acid. – HClO3: chloric acid – HClO4: perchloric acid Atoms, Molecules, and Ions •© 2009, Prentice-Hall, 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) . Atoms, Molecules, and Ions •© 2009, Prentice-Hall, Nomenclature of Binary Compounds • The ending on the more electronegative element is changed to -ide. – CO2: carbon dioxide – CCl4: carbon tetrachloride Atoms, Molecules, and Ions •© 2009, Prentice-Hall, 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. N2O5: dinitrogen pentoxide Atoms, Molecules, and Ions •© 2009, Prentice-Hall,