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Chapter 3 Molecules, Compounds, and Chemical Equations 1 Elements and Compounds • Elements combine together to make an almost limitless number of compounds. • The properties of the compound are totally different from the constituent elements. • Compounds are made of atoms held together by chemical bonds. • Bonds are forces of attraction between atoms. • The bonding attraction comes from attractions between protons and electrons 2 Bond Types • Two general types of bonding between atoms are found in compounds, ionic and covalent. • Ionic bonds result when electrons have been transferred between atoms, resulting in oppositely charged ions that attract each other. 9 generally found when metal atoms bonded to nonmetal atoms Usually metal + nonmetal • Covalent bonds result when two atoms share some of their electrons. 9 generally found when nonmetal atoms bonded together 3 3 1 Which of the following compounds exhibits both ionic and covalent bonding? a. b. c. d. e. SO2 CF4 NaCl Na2SO4 P4O10 4 5 5 Types of Formula: Empirical Formula • Empirical Formulas describe the kinds of elements found in the compound and the ratio of their atoms. 9They do not describe how many atoms, the order of attachment, or the shape. 9The formulas for ionic compounds are empirical. The empirical formula for the ionic compound fluorspar is CaCl2. This means that there is 1 Ca2+ ion for every 2 Cl− ions in the compound. The empirical formula for the molecular compound oxalic acid is CHO2. This means that there is 1 C atom and 1 H atom for every 2 O atoms in the molecule. The actual molecular formula is C2H2O46. 2 Types of Formula: Molecular Formula • Molecular Formulas describe the kinds of elements found in the compound and the exact numbers of their atoms. The molecular formula for oxalic acid is C2H2O4. This does not tell you that the shape of the molecule. 7 Types of Formula: Structural Formula • Structural Formulas describe the kinds of elements found in the compound, the numbers of their atoms, the order of atom attachment, and the kind of attachment. 9 use lines to represent covalent bonds 9 Each line describes the number of electrons shared by the bonded atoms. ¾ single line = 2 shared electrons, a single covalent bond ¾ double line = 4 shared electrons, a double covalent bond ¾ triple line = 6 shared electrons, a triple covalent bond 8 Example of all three types of structures Structural Formula of Oxalic Acid but Molecular Formula of Oxalic Acid C2H2O4. And Empirical Formula of Oxalic Acid (simplest ratio) CHO2. 9 3 Clicker question #3-1 H ethylene glycol C O What is the empirical formula for ethylene glycol? a) C2H6O2 b) CHO c) CH3O d) C2H3O2 10 Molecular View of Elements and Compounds 11 “Molecular” Elements • Certain elements occur as two-atom molecules. 9 Rule of 7’s • Other elements occur as polyatomic molecules. 9P4, S8, Se8 H2 7A N2 7 O2 F2 Cl2 Br2 I2 12 4 Ionic vs. Molecular Compounds Propane—contains individual C3H8 molecules Table salt—contains an array of Na+ ions and Cl– ions 13 Ionic Compounds • Compounds of metals with nonmetals are made of ions. 9Metal atoms form cations; nonmetal atoms form anions. • no individual molecule units, instead they have a three-dimensional array of cations and anions made of formula units • many contain polyatomic ions 9several atoms attached together in one ion examples; CO3-2 NO314 Write the formula of a compound made from aluminum ions and oxide ions. 1. Write the symbol for the +3 metal cation and its charge. Al column 3A 2. Write the symbol for the 3. 4. 5. O-2 column 6A nonmetal anion and its charge. Charge (without sign) becomes subscript for other ion. Reduce subscripts to smallest whole number ratio. Check that the total charge of the cations cancels the total charge of the anions. from Chapter 2 Al+3 O2Al2 O3 Al = (2) × (+3) = +6 O = (3) × (–2) = –6 15 5 Practice—What are the formulas for compounds made from the following ions? potassium ion with a nitride ion • K+ with N3– calcium ion with a bromide ion 16 Clicker question #3-2 What is the formula for the compound made from the following ions? Cr 6+ and O 21. 2. 3. 4. CrO Cr6O2 Cr2O6 CrO3 17 *Important Formula-to-Name *Rules for Ionic Compounds* • made of cation and anion • Write systematic name by simply naming the ions. 9if cation is: ¾metal with invariant charge = metal name ¾metal with variable charge = metal name (charge) ¾polyatomic ion = name of polyatomic ion 9if anion is: ¾nonmetal = stem of nonmetal name + -ide ¾polyatomic ion = name of polyatomic ion 18 6 Metal Cations 1) metals with invariant charge 9metals whose ions can have only one possible charge ¾Groups 1A & 2A , and Al, Ag, Zn, Sc 9cation name = metal name 19 2) Metals with variable charges metals whose ions can have more than one possible charge determine charge of cation by charge on anion name = metal name with Roman numeral charge in parentheses 20 Naming Nonmetal Anion • Determine the charge from position on the Periodic Table. • To name anion, change ending on the element name to -ide. 4A = 4− 5A = 3− 6A = 2− 7A = 1− C = carbide N = nitride O = oxide F = fluoride Si = silicide P = phosphide S = sulfide Cl = chloride Br = bromide I = iodide 21 7 REVIEW Predictable Ion Charges (for group A only) C4 Si4 - P3- 22 22 Naming Binary Ionic Compounds for Metals with Invariant Charge Name CsF 1. Identify cation and anion. Cs = Cs+ because it is in Group 1A F = F– because it is in Group 7A 2. Name the cation Cs+ = cesium 3. Name the anion F– = fluoride 4. Write the cation name first, then the anion name. 23 Cesium fluoride cesium fluoride Name the following compounds. 1. KCl 2. MgBr2 3. Al2S3 potassium chloride 24 8 Naming Binary Ionic Compounds for Metals with Variable Charge 1. Metal cation name is metal name followed by a Roman numeral in parentheses to indicate its charge. 9 determine charge from anion charge 9 common ions Table 3.4 2. nonmetal anion named by changing the ending on the nonmetal name to -ide 25 Find the charge on the cation. 1. TiCl4 2. CrO3 4 Cl = 4−, ∴ Ti = 4+ 26 Clicker question #3-3 Find the charge on Fe in the compound Fe3N2 1. 2. 3. 4. 5. +1 +2 +3 -3 +6 27 9 Example—Naming Binary Ionic with Variable Charge Metal CuF2 1. Identify cation and anion. F = F− because it is Group 7 Cu = Cu2+ to balance the two (−) charges from 2 F− 2. Name the cation. Cu2+ = copper(II) 3. Name the anion. F− = fluoride 4. Write the cation name first, then the anion name. copper(II) Copper (II)fluoride fluoride 28 Name the following compounds. 1. TiCl4 2. PbBr2 3. Fe2S3 titanium(IV) chloride 29 Practice—What are the formulas for compounds made from the following ions? copper(II) ion with a nitride ion iron(III) ion with a bromide ion 30 10 Compounds Containing Polyatomic Ions • Polyatomic ions are single ions that contain more than one atom. • Name and charge of polyatomic ion need to be learned ! 31 Some Common Polyatomic Ions Name Formula Name Formula acetate carbonate hydrogen carbonate (aka bicarbonate) hydroxide nitrate nitrite chromate dichromate ammonium C2H3O2– CO32– ClO– ClO2– ClO3– ClO4– SO42– SO32– HCO3– OH– NO3– NO2– CrO42– Cr2O72– NH4+ hypochlorite chlorite chlorate perchlorate sulfate sulfite hydrogen sulfate (aka bisulfate) hydrogen sulfite (aka bisulfite) phosphate HSO4– HSO3– PO43- 32 Special Pattern for Oxygenated halogens • -ate ion 9chlorate = ClO3– • -ate ion + one O ⇒ same charge, per- prefix 9perchlorate = ClO4– • -ate ion – one O ⇒ same charge, -ite suffix 9chlorite = ClO2– • -ate ion – two O ⇒ same charge, hypoprefix, -ite suffix 9hypochlorite = ClO– 33 11 Example—Naming Ionic Compounds Containing a Polyatomic Ion Fe(NO3)3 1. Identify the ions. 2. 3. 4. NO3 = NO3− a polyatomic ion Fe = Fe3+ to balance the charge of the 3 NO3− Name the cation. Fe3+ = iron(III), metal with variable charge Name the anion. NO3− = nitrate Write the name of the cation followed by the name of the anion. iron(III) Iron (III) nitrate 34 Name the following compounds 1. NH4Cl 2. Ca(ClO2)2 3. Cu(NO3)2 NH4+ Cl- ammonium chloride 35 Practice—What are the formulas for compounds made from the following ions? aluminum ion with a sulfate ion Al3+ with SO42− Al2(SO4)3 chromium(II) with hydrogen carbonate 36 12 Clicker question #3-4 All the following ions have the same charge ? EXCEPT 1) Sulfide ion 2) Bromate ion 3) Oxide ion 4) Sulfate ion 5) monohydrogen phosphate ion 37 Hydrates • Hydrates are ionic compounds • containing a specific number of waters for each formula unit. in formula, attached waters follow 9 CoCl2·6H2O • in name, attached waters indicated by prefix+hydrate after name of ionic compound 9 CoCl2·6H2O = cobalt(II) chloride hexahydrate 9 CaSO4·½H2O = calcium sulfate hemihydrate Prefix No. of Waters hemi ½ mono 1 di 2 tri 3 tetra 4 penta 5 hexa 6 hepta 7 octa 8 38 Writing Names of Binary Molecular Compounds (nonmetal + nonmetal) 1. Write the name of the first element in the formula. 9 element furthest left and down on the Periodic Table 9 Use the full name of the element, 2. Write the name of the second element in the formula with an -ide suffix. 3. Use a prefix in front of each name to indicate the number of atoms. a) Never use the prefix mono- on the first element. 39 13 Name the following. NO2 nitrogen dioxide PCl5 I2F7 40 Write formulas for the following. dinitrogen tetroxide N2O4 sulfur hexafluoride diarsenic trisulfide 41 Acids • Acids are molecular compounds that form H+ when dissolved in water. 9 To indicate the compound is dissolved in water, (aq) is written after the formula. HCl (aq) ¾ NOTE: not named as acid if not dissolved in water • dissolve many metals 9 like Zn, Fe, Mg; but not Au, Ag, Pt • formula generally starts with H 9 e.g., HBr (aq), H2SO4 (aq) 42 14 Acids • contain H+ cation and anion 9in aqueous solution • Binary acids have H+ cation and nonmetal anion. or HCl (aq) H2SO4(aq) •Oxyacids have H+ cation and polyatomic anion. 43 Naming Binary Acids • • • • Write a hydro prefix. Follow with the nonmetal name. Change ending on nonmetal name to -ic. Write the word acid at the end of the name. 44 Example—Naming Binary Acids: HCl(aq) 1. Identify the anion. Cl = Cl−, chloride because Group 7A 2. Name the anion with an -ic suffix. Cl− = chloride ⇒ chloric 3. Add a hydro- prefix to the anion name. hydrochloric 4. Add the word acid to the end. hydrochloric acid 45 15 Naming Oxyacids • If the polyatomic ion name ends in -ate, then change ending to -ic suffix. • If the polyatomic ion name ends in -ite, then change ending to -ous suffix. • Write word acid at the end of all names. 46 Example—Naming Oxyacids H2SO4(aq) 1. Identify the anion. SO4 = SO42− = sulfate 2. If the anion has -ate suffix, change it to -ic. If the anion has -ite suffix, change it to -ous. SO42− = sulfate ⇒ sulfuric 3. Write the name of the anion followed by the word acid. sulfuric Sulfuricacid acid 47 Name the following H2S (aq) hydrosulfuric acid HClO3 (aq) HNO2 (aq) 48 16 Clicker question #3-5 Choose the name-formula pair that does NOT match. 1) ammonium hydrogen carbonate, (NH4)2CO3 2) sodium chlorite, NaClO2 3) calcium hydride, CaH2 4) nitric acid, HNO3 5) calcium hydroxide, Ca(OH)2 49 Writing Formulas for Acids • when name ends in acid, formula starts with H • Write formulas as if ionic, even though it is molecular. • Hydro prefix means it is a binary acid, no hydro prefix means it is an oxyacid. • for oxyacid, if ending is -ic, polyatomic ion ends in -ate; – if ending is -ous, polyatomic ion ends in -ite 50 Practice—What are the formulas for the following acids? chlorous acid H+ with ClO2– HClO2 phosphoric acid 51 17 Formula Mass • is the mass of an individual molecule or formula unit 9also known as molecular mass or molecular weight • sum of the masses of the atoms in a single molecule or formula unit mass of 1 molecule of H2O = 2(1.01 amu H) + 16.00 amu O = 18.02 amu 52 Molar Mass of Compounds • The masses of molecules can be calculated from atomic masses. • since 1 mole of H2O contains 2 moles of H and 1 mole of O Molar Mass = 1 mole H2O = 2(1.01 g H) +16.00 g O = 18.02 g Molar Mass of H2O is 18.02 g/mole NOTE! You will need to more precise molar masses during tests and assignments. A more precise molar mass of water uses the exact atomic weight from the peroidic table. Molar Mass = 2(1.0079 g H) +15.9999 g O = 18.0148 g/mole of H2O 53 Practice—How many moles are in 50.0 g of PbO2? Given: 50.0 g PbO2 Find: moles PbO2 Conceptual Plan: g PbO2 mol PbO2 Relationships: 1 mol PbO2 = 239.2 g Solution: Check: Since the given amount is less than 239.2 g, the moles being <1 makes sense. 54 18 Example Find the number of CO2 molecules in 10.8 g of dry ice. Given: 10.8 g CO2 Find: molecules CO2 g CO2 Conceptual Plan: Relationships: mol CO2 molec CO2 1 mol CO2 = 44.01 g, 1 mol = 6.022 × 1023 Solution: Check: Since the given amount is much less than 1 mol CO2, the number makes sense. 55 Clicker question #3-6 How many molecules are in 1.00 g of mannose, C 6H 12 O 6, which is a sweet tasting sugar that has a bitter aftertaste? 1) 2.18 x 10 18 2) 3.34 x 10 18 3) 2.18 X 10 21 4) 3.34 x 10 21 5) 3.34 x 10 24 56 Percent Composition • percentage of each element in a compound 9 by mass • can be determined from 1. the formula of the compound 2. the experimental mass analysis of the compound 57 19 Example 3.13 Find the mass percent of Cl in C2Cl4F2. Given: Find: C2Cl4F2 % Cl by mass Conceptual Plan: Relationships: Solution: Check: Since the percentage is less than 100 and Cl is much heavier than the other atoms, the number makes sense. 58 Practice—Determine the mass percent composition of the CaCl2 59 Clicker question #3-7 The mineral leadhillite, which is essentially Pb4(SO4)(CO3)2(OH)2 (FW = 1079 amu), contains ____ % oxygen by weight. 1) 10.4 2) 11.9 3) 13.3 4) 14.8 5) 17.8 60 20 Empirical Formula • simplest, whole-number ratio of the atoms of elements in a compound • can be determined from elemental analysis by: method 1. percent composition or Method 2. combustion analysis i.e., masses of elements formed from decomposition or reaction 61 1. Empirical Formula from % mass 1) Convert the percentages to grams. - assume you start with 100.00 g of the compound 2) Convert grams to moles. - use molar mass of each element 3) Write a pseudo-formula using moles as subscripts 4) Divide all by smallest number of moles. - If result is within 0.1 of whole number, round to whole number. 5) If required, multiply all mole ratios by number to make all whole numbers. - if ratio 0.5, multiply all by 2; if ratio 0.33 or 0.67, multiply all by 3; if ratio 0.25 or 0.75, multiply all by 4; etc. 62 Example 3.16 • Laboratory analysis of aspirin determined the following mass percent composition. Find the empirical formula. C = 60.00% H = 4.48% O = 35.53% 63 21 Strategy The empirical formula is CxHyOz and found by: ggCC molCC mol ggHH molHH mol O ggO molO O mol mole pseudo- ratio pseudoformula formula whole number ratio empirical empirical formula formula Masses found in step #1 (next slide) Step 1 Assume 100 g of sample; therefore there are 60.00 g C, 4.48 g H, and 35.53 g O. Step 2 Find moles Step 3 write a pseudoformula C4.996H4.44O2.220 find the mole ratio by dividing by the smallest number of moles Step 4 multiply subscripts by factor to give whole number {C2.25H2O1} x 4 =C9H8O4 22 Clicker question #3-8 Analysis of a hydrocarbon showed that it contained 14.4% hydrogen and 85.6% carbon by weight. What is its simplest formula? 1) CH 2) CH 2 3) CH 3 4) C 2H 3 5) C 2H 5 67 Molecular Formulas • • The molecular formula is a multiple of the empirical formula. To determine the molecular formula, you need to know the empirical formula and the molar mass of the compound. 68 68 Example 3.17 Find the molecular formula of butanedione given the emp. formula = C2H3O and the molar mass is 86.09g /mole Check: The molar mass of the calculated formula is in agreement with the given molar mass. 69 23 Empirical formula from Combustion Analysis • A common technique for analyzing compounds is to burn a known mass of compound and weigh the amounts of product made. 9 generally used for organic compounds containing C, H, O • By knowing the mass of the product and composition of constituent elements in the product, the original amount of constituent elements can be determined. 9 All the original C forms CO2, the original H forms H2O, the original mass of O is found by subtraction. • Once the masses of all the constituent elements in the original compound have been determined, the empirical formula can be found. 70 71 Example 3.19 • Combustion of a 0.8233-g sample of a compound containing only carbon, hydrogen, and oxygen produced the following: CO2 = 2.445 g H2O = 0.6003 g Determine the empirical formula of the compound. 72 24 Write a conceptual plan for example 3.19 g CO2, H2O mol C, H, O mol CO2, H2O mol C, H pseudo formula g C, H mol ratio g O mol O empirical formula Example 3.19 (continued) Step 1 Calculate the moles of C and H. Step 2 Calculate the grams of C and H Step 3 Calculate the grams and moles of O. Step 4 Write a pseudoformula using moles found of each element C0.05556H0.06662O0.00556 divide by the smallest number of moles. 75 25 Practice 1 A 47.95 mg sample of a compound containing only C and H was combusted and produced 145.21 mg of CO2 and 74.33 mg of H2O. (a) What is the empirical formula of this hydrocarbon compound? (b) The molar mass of the hydrocarbon was found to be 87.18 g mol-1. What is its molecular formula? Answer is: a) C2H5 b) C6H15 76 Clicker question #3-9 A given hydrocarbon is converted completely to carbon dioxide and water, and equal numbers of moles of CO2 and H2O are produced. The hydrocarbon could be 1) CH4. 2) C2H6. 3) C3H8. 4) C3H4. 5) C4H8. 77 Chemical Reactions • Reactions involve chemical changes in matter resulting in new substances. • Reactions involve rearrangement and exchange of atoms to produce new molecules. Reactants → Products 78 78 26 Chemical Equations • shorthand way of describing a reaction • provides information about the reaction 9formulas of reactants and products 9states of reactants and products 9relative numbers of reactant and product molecules that are required 9can be used to determine weights of reactants used and products that can be made 79 Combustion of Methane CH4(g) + O2(g) → CO2(g) + H2O(g) (unbalanced) • To show the reaction obeys the Law of Conservation of Mass, the equation must be balanced. 9 We adjust the numbers of molecules so there are equal numbers of atoms of each element on both sides of the arrow. CH4(g) + 2 O2(g) → CO2(g) + 2 H2O(g) Check equal # of elements 1C + 4H + 4O 1C + 4H + 4O 80 Balancing Chemical Equations CH4(g) + 2 O2(g) → CO2(g) + 2 H2O(g) • CH4 and O2 are the reactants, and CO2 and H2O are the products. • The (g) after the formulas tells us the state of the chemical. • The number in front of each substance tells us the number of those molecules in the reaction. 9 called the coefficients 81 27 Symbols Used in Equations • symbols used to indicate state after chemical 9(g) = gas; (l) = liquid; (s) = solid 9(aq) = aqueous = dissolved in water • energy symbols used above the arrow for decomposition reactions 9 Δ = heat 9 hν = light ΔH (chapter 6) 82 Example 3.21 Write a balanced equation for the combustion of butane, C4H10. Write a skeletal equation: Balance atoms in complex substances first: Balance free elements by adjusting coefficient in front of free element: C4H10(l) + O2(g) → CO2(g) + H2O(g) 4⇐C⇒1×4 C4H10(l) + O2(g) → 4 CO2(g) + H2O(g) 10 ⇐ H ⇒ 2 × 5 C4H10(l) + O2(g) → 4 CO2(g) + 5 H2O(g) 13/2 × 2 ⇐ O ⇒ 13 C4H10(l) + 13/2 O2(g) → 4 CO2(g) + 5 H2O(g) If fractional coefficients, {C4H10(l) + 13/2 O2(g) → 4 CO2(g) + 5 H2O(g)} × 2 multiply thru by 2 C4H10(l) + 13 O2(g) → 8 CO2(g) + 10 H2O(g) denominator: Check: 8 ⇐ C ⇒ 8; 20 ⇐ H ⇒ 20; 26 ⇐ O ⇒ 26 83 Practice 1 When aluminum metal reacts with air, it produces a white, powdery compound— aluminum oxide, Al2O3. Write a balanced equation including proper physical states. 84 28 Clicker question #3-10 Which of the following equations is (are) balanced? → PbCl2 + NaNO 3 a. NaCl + Pb(NO 3)2 ⎯ ⎯ → N 2 + 4H 2O + Cr2O 3 b. (NH 4)2Cr2O 7 ⎯ ⎯ → 3Fe + 3CO 2 c. Fe3O 4 + 3CO ⎯ ⎯ 1) a only 2) b only 3) c only 4) a and b only 5) a, b, and c 85 29