Presentation on theme: "Chapter 7 Chemical Quantities"— Presentation transcript:
1Chapter 7 Chemical Quantities Charles Page High SchoolDr. Stephen L. Cotton
2Section 7.1 The Mole: A Measurement of Matter OBJECTIVES:Describe how Avogadro’s number is related to a mole of any substance.
3Section 7.1 The Mole: A Measurement of Matter OBJECTIVES:Calculate the mass of a mole of any substance.
4What is a Mole? You can measure mass, or volume, or you can count pieces.We measure mass in grams.We measure volume in liters.We count pieces in MOLES.
5Moles (abbreviated: mol) Defined as the number of carbon atoms in exactly 12 grams of carbon-12.1 mole is 6.02 x particles.Treat it like a very large dozen6.02 x is called Avogadro’s number.
6Representative particles The smallest pieces of a substance.For a molecular compound: it is the molecule.For an ionic compound: it is the formula unit (ions).For an element: it is the atom.Remember the 7 diatomic elements (made of molecules)
7Types of questions How many oxygen atoms in the following? CaCO3 Al2(SO4)3How many ions in the following?CaCl2NaOH
8Types of questionsHow many molecules of CO2 are there in 4.56 moles of CO2 ?How many moles of water is 5.87 x molecules?How many atoms of carbon are there in 1.23 moles of C6H12O6 ?How many moles is 7.78 x 1024 formula units of MgCl2?
9Measuring Moles Remember relative atomic mass? The amu was one twelfth the mass of a carbon-12 atom.Since the mole is the number of atoms in 12 grams of carbon-12,the decimal number on the periodic table is also the mass of 1 mole of those atoms in grams.
10Gram Atomic Mass (gam)Equals the mass of 1 mole of an element in grams12.01 grams of C has the same number of pieces as grams of H and grams of iron.We can write this as g C = 1 mole CWe can count things by weighing them.
11Examples How much would 2.34 moles of carbon weigh? How many moles of magnesium is g of Mg?How many atoms of lithium is 1.00 g of Li?How much would 3.45 x 1022 atoms of U weigh?
12What about compounds?in 1 mole of H2O molecules there are two moles of H atoms and 1 mole of O atomsTo find the mass of one mole of a compounddetermine the moles of the elements they haveFind out how much they would weighadd them up
13What about compounds? What is the mass of one mole of CH4? 1 mole of C = g4 mole of H x 1.01 g = 4.04g1 mole CH4 = = 16.05gThe Gram Molecular Mass (gmm) of CH4 is 16.05gthis is the mass of one mole of a molecular compound.
14Gram Formula Mass (gfm) The mass of one mole of an ionic compound.Calculated the same way as gmm.What is the GFM of Fe2O3?2 moles of Fe x g = g3 moles of O x g = gThe GFM = g g = g
15Section 7.2 Mole-Mass and Mole-Volume Relationships OBJECTIVES:Use the molar mass to convert between mass and moles of a substance.
16Section 7.2 Mole-Mass and Mole-Volume Relationships OBJECTIVES:Use the mole to convert among measurements of mass, volume, and number of particles.
17Molar MassMolar mass is the generic term for the mass of one mole of any substance (in grams)The same as: 1) gram molecular mass, 2) gram formula mass, and 3) gram atomic mass- just a much broader term.
18ExamplesCalculate the molar mass of the following and tell what type it is:Na2SN2O4CCa(NO3)2C6H12O6(NH4)3PO4
19Molar Mass The number of grams of 1 mole of atoms, ions, or molecules. We can make conversion factors from these.To change grams of a compound to moles of a compound.
22For example How many moles is 5.69 g of NaOH? need to change grams to moles
23For example How many moles is 5.69 g of NaOH? need to change grams to molesfor NaOH
24For example How many moles is 5.69 g of NaOH? need to change grams to molesfor NaOH1mole Na = 22.99g 1 mol O = g 1 mole of H = 1.01 g
25For example How many moles is 5.69 g of NaOH? need to change grams to molesfor NaOH1mole Na = 22.99g 1 mol O = g 1 mole of H = 1.01 g1 mole NaOH = g
26For example How many moles is 5.69 g of NaOH? need to change grams to molesfor NaOH1mole Na = 22.99g 1 mol O = g 1 mole of H = 1.01 g1 mole NaOH = g
27For example How many moles is 5.69 g of NaOH? need to change grams to molesfor NaOH1mole Na = 22.99g 1 mol O = g 1 mole of H = 1.01 g1 mole NaOH = g
28Examples How many moles is 4.56 g of CO2? How many grams is 9.87 moles of H2O?How many molecules is 6.8 g of CH4?49 molecules of C6H12O6 weighs how much?
29Gases Many of the chemicals we deal with are gases. They are difficult to weigh.Need to know how many moles of gas we have.Two things effect the volume of a gasTemperature and pressureWe need to compare them at the same temperature and pressure.
30Standard Temperature and Pressure 0ºC and 1 atm pressureabbreviated STPAt STP 1 mole of gas occupies 22.4 LCalled the molar volume1 mole = 22.4 L of any gas at STP
31Examples What is the volume of 4.59 mole of CO2 gas at STP? How many moles is L of O2 at STP?What is the volume of 8.8 g of CH4 gas at STP?
32Density of a gas D = m / V for a gas the units will be g / L We can determine the density of any gas at STP if we know its formula.To find the density we need the mass and the volume.If you assume you have 1 mole, then the mass is the molar mass (from PT)At STP the volume is 22.4 L.
33Examples Find the density of CO2 at STP. Find the density of CH4 at STP.
34The other wayGiven the density, we can find the molar mass of the gas.Again, pretend you have 1 mole at STP, so V = 22.4 L.m = D x Vm is the mass of 1 mole, since you have 22.4 L of the stuff.What is the molar mass of a gas with a density of g/L?2.86 g/L?
35Summary These four items are all equal: a) 1 mole b) molar mass (in grams)c) 6.02 x 1023 representative particlesd) 22.4 L at STPThus, we can make conversion factors from them.
36Section 7.3 Percent Composition and Chemical Formulas OBJECTIVES:Calculate the percent composition of a substance from its chemical formula or experimental data.
37Section 7.3 Percent Composition and Chemical Formulas OBJECTIVES:Derive the empirical formula and the molecular formula of a compound from experimental data.
38Calculating Percent Composition of a Compound Like all percent problems:Part wholeFind the mass of each component,then divide by the total mass.x 100 %
39ExampleCalculate the percent composition of a compound that is 29.0 g of Ag with 4.30 g of S.
40Getting it from the formula If we know the formula, assume you have 1 mole.Then you know the mass of the pieces and the whole.
41Examples Calculate the percent composittion of C2H4? How about Aluminum carbonate?Sample Problem 7-11, p.191We can also use the percent as a conversion factorSample Problem 7-12, p.191
42The Empirical FormulaThe lowest whole number ratio of elements in a compound.The molecular formula = the actual ratio of elements in a compound.The two can be the same.CH2 is an empirical formulaC2H4 is a molecular formulaC3H6 is a molecular formulaH2O is both empirical & molecular
43Calculating Empirical Just find the lowest whole number ratioC6H12O6CH4NIt is not just the ratio of atoms, it is also the ratio of moles of atoms.In 1 mole of CO2 there is 1 mole of carbon and 2 moles of oxygen.In one molecule of CO2 there is 1 atom of C and 2 atoms of O.
44Calculating Empirical We can get a ratio from the percent composition.Assume you have a 100 g.The percentages become grams.Convert grams to moles.Find lowest whole number ratio by dividing by the smallest.
45ExampleCalculate the empirical formula of a compound composed of % C, % H, and %N.Assume 100 g so38.67 g C x 1mol C = mole C gC16.22 g H x 1mol H = mole H gH45.11 g N x 1mol N = mole N gN
46Example The ratio is 3.220 mol C = 1 mol C 3.219 molN 1 mol N The ratio is mol H = 5 mol H molN mol N= C1H5N1A compound is % P and % O. What is the empirical formula?Caffeine is 49.48% C, 5.15% H, 28.87% N and 16.49% O. What is its empirical formula?
47Empirical to molecular Since the empirical formula is the lowest ratio, the actual molecule would weigh more.By a whole number multiple.Divide the actual molar mass by the empirical formula mass.Caffeine has a molar mass of 194 g. what is its molecular formula?
48ExampleA compound is known to be composed of % Cl, 24.27% C and 4.07% H. Its molar mass is known (from gas density) to be g. What is its molecular formula?Sample Problem 7-14, p.194