Physical and Chemical Changes and Properties
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1 Physical and Chemical Changes and Properties Practice 1.1 pg. 11 Determine whether each change is physical or chemical. What kind of property (chemical or physical) is demonstrated in each case? a. A copper wire is hammered flat. b. A nickel dissolves in acid to form a blue-green solution. c. Dry ice sublimes without melting. d. A match ignites when struck on a flint.
2 Converting between Temperature Scales For Practice 1.2 pg. 16 Gallium is a solid metal at room temperature but will melt to a liquid in your hand. The melting point of gallium is 85.6 F. What is this temperature on (a) the Celsius scale and (b) the Kelvin scale?
3 Calculating Density For Practice 1.3 pg. 19 The woman in this example is shocked that the ring is fake and returns it. She buys a new ring that has a mass of 4.53 g and a volume of cm 3. Is the new ring genuine? For More Practice 1.3 A metal cube has an edge length of 11.4 mm and a mass of 6.67 g. Calculate the density of the metal and use Table 1.4 to determine the likely identity of the metal.
4 Reporting the Correct Number of Digits For Practice 1.4 pg. 21 Record the temperature on the thermometer shown at the right to the correct number of digits.
5 Determining the Number of Significant Figures in a Number For Practice 1.5 pg. 23 How many significant figures are in each number? a. 554 km b. 7 pennies c m d s e km f. 21,000 m
6 Significant Figures in Calculations For Practice 1.6 pg. 25 Perform each calculation to the correct number of significant figures. a b. c ( ) d. ( )
7 Unit Conversion For Practice 1.7 pg. 29 Convert 288 cm to yards.
8 Unit Conversion For Practice 1.8 pg. 29 Convert 9255 cm 3 to gallons.
9 Unit Conversions Involving Units Raised to a Power For Practice 1.9 pg. 30 How many cubic centimeters are there in 2.11 yd 3? For More Practice 1.9 A vineyard has 145 acres of Chardonnay grapes. A particular soil supplement requires 5.50 grams for every square meter of vineyard. How many kilograms of the soil supplement are required for the entire vineyard? (1 km 2 = 247 acres)
10 Density as a Conversion Factor.For Practice 1.10 pg. 31 Backpackers often use canisters of white gas to fuel a cooking stove s burner. If one canister contains 1.45 L of white gas, and the density of the gas is g/cm 3, what is the mass of the fuel in kilograms? For More Practice 1.10 A drop of gasoline has a mass of 22 mg and a density of g/cm 3. What is its volume in cubic centimeters?
11 Problems with Equations For Practice 1.11 pg. 33 Find the radius (r) of an aluminum cylinder that is 2.00 cm long and has a mass of 12.4 g. For a cylinder, V = πr 2 l.
12 Problems with Equations For Practice 1.12 pg. 33 Find the density, in g/cm 3, of a metal cube with a mass of 50.3 g and an edge length (l) of 2.65 cm. For a cube, V = l 3.
13 Example 2.1 Law of Definite Proportions For Practice 2.1 Two samples of carbon monoxide are decomposed into their constituent elements. One sample produces 17.2 g of oxygen and 12.9 g of carbon, and the other sample produces 10.5 g of oxygen and 7.88 g of carbon. Show that these results are consistent with the law of definite proportions.
14 Law of Multiple Proportions For Practice 2.2 Hydrogen and oxygen form both water and hydrogen peroxide. The decomposition of a sample of water forms g hydrogen to every 1.00 g oxygen. The decomposition of a sample of hydrogen peroxide forms g hydrogen to every 1.00 g oxygen. Show that these results are consistent with the law of multiple proportions.
15 Example 2.3 Atomic Numbers, Mass Numbers, and Isotope Symbols Continued For Practice 2.3 a. What are the atomic number, mass number, and symbol for the carbon isotope with seven neutrons? b. How many protons and neutrons are present in an atom of?
16 Example 2.4 Predicting the Charge of Ions Predict the charges of the monoatomic (single atom) ions formed by these main-group elements. a. Al b. S Solution a. Aluminum is a main-group metal and tends to lose electrons to form a cation with the same number of electrons as the nearest noble gas. Aluminum atoms have 13 electrons and the nearest noble gas is neon, which has 10 electrons. Aluminum therefore loses 3 electrons to form a cation with a 3+ charge (Al 3+ ). b. Sulfur is a nonmetal and tends to gain electrons to form an anion with the same number of electrons as the nearest noble gas. Sulfur atoms have 16 electrons and the nearest noble gas is argon, which has 18 electrons. Sulfur therefore gains 2 electrons to form an anion with a 2 charge (S 2 ). For Practice 2.4 a. N b. Rb
17 Example 2.5 Atomic Mass Copper has two naturally occurring isotopes: Cu-63 with a mass of amu and a natural abundance of 69.17%, and Cu-65 with a mass of amu and a natural abundance of 30.83%. Calculate the atomic mass of copper. Solution Convert the percent natural abundances into decimal form by dividing by 100. Calculate the atomic mass using the equation given in the text. For Practice 2.5 Magnesium has three naturally occurring isotopes with masses of amu, amu, and amu and natural abundances of 78.99%, 10.00%, and 11.01%, respectively. Calculate the atomic mass of magnesium. For More Practice 2.5 Gallium has two naturally occurring isotopes: Ga-69 with a mass of amu and a natural abundance of 60.11%, and Ga-71. Use the atomic mass of gallium from the periodic table to find the mass of Ga-71.
18 Example 2.6 Converting between Number of Moles and Number of Atoms Continued Solution Check Since atoms are small, it makes sense that the answer is large. The given number of moles of copper is almost 2.5, so the number of atoms is almost 2.5 times Avogadro s number. For Practice 2.6 A pure silver ring contains silver atoms. How many moles of silver atoms does it contain?
19 Example 2.7 Converting between Mass and Amount (Number of Moles) Continued Solution Check The given mass of carbon is much less than the molar mass of carbon, so it makes sense that the answer (the amount in moles) is much less than 1 mol of carbon. For Practice 2.7 Calculate the amount of copper (in moles) in a 35.8 g pure copper sheet. For More Practice 2.7 Calculate the mass (in grams) of mol of titanium.
20 Example 2.8 The Mole Concept Converting between Mass and Number of Atoms Continued Solve Follow the conceptual plan to solve the problem. Begin with 3.10 g Cu and multiply by the appropriate conversion factors to arrive at the number of Cu atoms. Solution Check The answer (the number of copper atoms) is less than (1 mole). This is consistent with the given mass of copper atoms, which is less than the molar mass of copper. For Practice 2.8 How many carbon atoms are there in a 1.3-carat diamond? Diamonds are a form of pure carbon. (1 carat = 0.20 grams) For More Practice 2.8 Calculate the mass of tungsten atoms.
21 Example 2.9 The Mole Concept Continued Solution Check The units of the answer (cm) are correct. The magnitude cannot be estimated accurately, but a radius of about one-half of a centimeter is reasonable for just over one-tenth of a mole of aluminum atoms. For Practice 2.9 A titanium cube contains atoms. What is the edge length of the cube? The density of titanium is 4.50 g/cm 3. For More Practice 2.9 Find the number of atoms in a copper rod with a length of 9.85 cm and a radius of 1.05 cm. The density of copper is 8.96 g/cm 3.
22 Example 3.1 Molecular and Empirical Formulas Boric acid, H 3 BO 3, is a weak electrolyte. Would you expect a boric acid solution to contain only ions, only molecules, or mostly molecules and a few ions? a. C 5 H 12 b. Hg 2 Cl 2 c. C 2 H 4 O 2
23 Classifying Substances as Atomic Elements, Molecular Elements, Molecular Compounds, or Ionic Compounds For Practice 3.2 Classify each of the substances as an atomic element, molecular element, molecular compound, or ionic compound. a. fluorine b. N 2 O c. silver d. K 2 O e. Fe 2 O 3 FIGURE 3.5 Molecular Elements The highlighted elements exist primarily as diatomic molecules (yellow) or polyatomic molecules (red).
24 Writing Formulas for Ionic Compounds For Practice 3.3 Write the formula for the compound formed between potassium and sulfur.
25 Example 3.4 Writing Formulas for Ionic Compounds Write the formula for the ionic compound that forms between calcium and oxygen. Step 2 Adjust the subscript on each cation and anion to balance the overall charge. Step 3 Check that the sum of the charges of the cations equals the sum of the charges of the anions. cations: 2+ anions: 2 The charges cancel. For Practice 3.4 Write the formula for the compound formed between aluminum and nitrogen.
26 Example 3.5 Naming Ionic Compounds Containing a Metal That Forms Only One Type of Cation Name the compound CaBr 2. Solution The cation is calcium. The anion is from bromine, which becomes bromide. The correct name is calcium bromide. For Practice 3.5 Name the compound Ag 3 N. For More Practice 3.5 Write the formula for rubidium sulfide.
27 Example 3.6 Naming Ionic Compounds Containing a Metal That Forms More Than One Kind of Cation Name the compound PbCl 4. Solution The charge on Pb must be 4+ for the compound to be charge-neutral with four Cl anions. The name for PbCl 4 is the name of the cation, lead, followed by the charge of the cation in parentheses (IV) and the base name of the anion, chlor, with the ending -ide. The full name is lead(iv) chloride. PbCl 4 lead(iv) chloride For Practice 3.6 Name the compound FeS. For More Practice 3.6 Write the formula for ruthenium(iv) oxide.
28 Example 3.7 Naming Ionic Compounds That Contain a Polyatomic Ion Name the compound Li 2 Cr 2 O 7. Solution The name for Li 2 Cr 2 O 7 is the name of the cation, lithium, followed by the name of the polyatomic ion, dichromate. Its full name is lithium dichromate. Li 2 Cr 2 O 7 lithium dichromate For Practice 3.7 Name the compound Sn(ClO 3 ) 2. For More Practice 3.7 Write the formula for cobalt(ii) phosphate.
29 Example 3.8 Naming Molecular Compounds Name each compound. a. NI 3 b. PCl 5 c. P 4 S 10 Solution a. The name of the compound is the name of the first element, nitrogen, followed by the base name of the second element, iod, prefixed by tri- to indicate three and given the suffix -ide. b. The name of the compound is the name of the first element, phosphorus, followed by the base name of the second element, chlor, prefixed by penta- to indicate five and given the suffix -ide. c. The name of the compound is the name of the first element, phosphorus, prefixed by tetra- to indicate four, followed by the base name of the second element, sulf, prefixed by deca to indicate ten and given the suffix -ide. For Practice 3.8 Name the compound N 2 O 5. For More Practice 3.8 Write the formula for phosphorus tribromide.
30 Example 3.9 Naming Binary Acids Name HI(aq). Solution The base name of I is iod, so HI(aq) is hydroiodic acid. HI(aq) hydroiodic acid For Practice 3.9 Name HI(aq).
31 Example 3.10 Naming Binary Oxyacids Name HC 2 H 3 O 2 (aq). Solution The oxyanion is acetate, which ends in -ate; therefore, the name of the acid is acetic acid. HC 2 H 3 O 2 (aq) acetic acid For Practice 3.10 Name HNO 2 (aq). For More Practice 3.10 Write the formula for perchloric acid.
32 Example 3.11 Using the Nomenclature Flow Chart to Name Compounds Continued c. CoF 2 Begin by determining whether the compound is ionic, molecular, or an acid. Since CoF 2 contains a metal and a nonmetal, it is ionic. Next refer to Table 3.2 to determine whether the metal forms one type of ion or more than one type. Since Co is not listed in Table 3.2, it must form more than one type of ion. Name the compound as the name of the cation, cobalt, followed by the charge of the cation in parentheses (II), and the base name of the anion, fluor, with the ending -ide. For Practice 3.11 Use the flow chart in Figure 3.10 to name H 2 SO 3 (aq).
33 Example 3.12 Calculating Formula Mass Calculate the formula mass of glucose, C 6 H 12 O 6. Solution To find the formula mass, add the atomic masses of each atom in the chemical formula: For Practice 3.12 Calculate the formula mass of calcium nitrate.
34 Example 3.13 The Mole Concept Converting between Mass and Number of Molecules Continued For Practice 3.13 Find the number of ibuprofen molecules in a tablet containing mg of ibuprofen (C 13 H 8 O 2 ). For More Practice 3.13 What is the mass of a sample of water containing H 2 O molecules?
35 Example 3.14 Mass Percent Composition Continued Solution Check The units of the answer (%) are correct, and the magnitude is reasonable because (a) it is between 0 and 100% and (b) chlorine is the heaviest atom in the molecule and there are four atoms of it. For Practice 3.14 Acetic acid (C 2 H 4 O 2 ) is the active ingredient in vinegar. Calculate the mass percent composition of oxygen in acetic acid. For More Practice 3.14 Calculate the mass percent composition of sodium in sodium oxide.
36 Example 3.15 Using Mass Percent Composition as a Conversion Factor Continued Solve Follow the conceptual plan to solve the problem. Solution You can consume 6.2 g NaCl and still be within the FDA guidelines. Check The units of the answer are correct. The magnitude seems reasonable because it is larger than the amount of sodium, as expected, because sodium is only one of the elements in NaCl. For Practice 3.15 What mass (in grams) of iron(iii) oxide contains 58.7 grams of iron? Iron(III) oxide is 69.94% iron by mass. For More Practice 3.15 If someone consumes 22 g of sodium chloride per day, what mass (in grams) of sodium does that person consume? Sodium chloride is 39% sodium by mass packets of salt contain 6.2 g of NaCl.
37 Example 3.16 Chemical Formulas as Conversion Factors Continued Solve Follow the conceptual plan to solve the problem. Solution Check The units of the answer (g H) are correct. Since a gallon of water is about 3.8 L, its mass is about 3.8 kg. H is a light atom, so its mass should be significantly less than 3.8 kg, as it is in the answer. For Practice 3.16 Determine the mass of oxygen in a 7.2 g sample of Al 2 (SO 4 ) 3. For More Practice 3.16 Butane (C 4 H 10 ) is the liquid fuel in lighters. How many grams of carbon are present within a lighter containing 7.25 ml of butane? (The density of liquid butane is g/ml.)
38 Example 3.17 Obtaining an Empirical Formula from Experimental Data Continued For Practice 3.17 A sample of a compound is decomposed in the laboratory and produces 165 g carbon, 27.8 g hydrogen, and g oxygen. Calculate the empirical formula of the compound.
39 Example 3.18 Obtaining an Empirical Formula from Experimental Data Continued For Practice 3.18 Ibuprofen has the following mass percent composition: C 75.69%, H 8.80%, O 15.51%. What is the empirical formula of ibuprofen?
40 Example 3.19 Calculating a Molecular Formula from an Empirical Formula and Molar Mass Continued Divide the molar mass by the empirical formula mass to find n. Multiply the empirical formula by n to obtain the molecular formula. Check Check the answer by calculating the molar mass of the formula as follows: 4(12.01 g/mol) + 6(1.008 g/mol) + 2(16.00 g/mol) = g/mol The calculated molar mass is in agreement with the given molar mass. For Practice 3.19 A compound has the empirical formula CH and a molar mass of g/mol. What is its molecular formula?
41 Example 3.19 Calculating a Molecular Formula from an Empirical Formula and Molar Mass Continued For More Practice 3.19 A compound with the percent composition shown next has a molar mass of g/mol. Determine its molecular formula. C, 39.97% H, 13.41% N, 46.62%
42 Example 3.20 Determining an Empirical Formula from Combustion Analysis Continued Step 6 Divide all the subscripts in the formula by the smallest subscript. (Round all subscripts that are within 0.1 of a whole number.) Step 7 If the subscripts are not whole numbers, multiply all the subscripts by a small whole number to get whole-number subscripts. C 1 H 24 C 5 H 24 The correct empirical formula is C 5 H 12. For Practice 3.20 Upon combustion, a compound containing only carbon and hydrogen produced 1.60 g CO 2 and g H 2 O. Find the empirical formula of the compound.
43 Example 3.21 Determining an Empirical Formula from Combustion Analysis Continued Step 5 Write down a pseudoformula for the compound using the number of moles of each element (from steps 3 and 4) as subscripts. C H O Step 6 Divide all the subscripts in the formula by the smallest subscript. (Round all subscripts that are within 0.1 of a whole number.) Step 7 If the subscripts are not whole numbers, multiply all the subscripts by a small whole number to get whole-number subscripts. The subscripts are whole numbers; no additional multiplication is needed. The correct empirical formula is C 10 H 12 O. For Practice 3.21 Upon combustion, a g sample of a compound containing only carbon, hydrogen, and oxygen produced g CO 2 and g H 2 O. Find the empirical formula of the compound.
44 Example 3.22 Balancing Chemical Equations Continued Step 5 Check to make certain the equation is balanced by summing the total number of each type of atom on both sides of the equation. The equation is balanced. For Practice 3.22 Write a balanced equation for the reaction between solid silicon dioxide and solid carbon to produce solid silicon carbide and carbon monoxide gas.
45 Example 3.23 Balancing Chemical Equations Continued Step 5 Check to make certain the equation is balanced by summing the total number of each type of atom on both sides of the equation. The equation is balanced. For Practice 3.23 Write a balanced equation for the combustion of gaseous ethane (C 2 H 6 ), a minority component of natural gas, in which it combines with gaseous oxygen to form gaseous carbon dioxide and gaseous water.
46 Example 3.24 Balancing Chemical Equations Containing Ionic Compounds with Polyatomic Ions Continued For Practice 3.24 Write a balanced equation for the reaction between aqueous lead(ii) nitrate and aqueous potassium chloride to form solid lead(ii) chloride and aqueous potassium nitrate.
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