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1 Chemical Quantities or. 2 How you measure how much? How you measure how much? n You can measure mass, n or volume, n or you can count pieces. n We measure.

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Presentation on theme: "1 Chemical Quantities or. 2 How you measure how much? How you measure how much? n You can measure mass, n or volume, n or you can count pieces. n We measure."— Presentation transcript:

1 1 Chemical Quantities or

2 2 How you measure how much? How you measure how much? n You can measure mass, n or volume, n or you can count pieces. n We measure mass in grams. n We measure volume in liters. n We count pieces in MOLES.

3 3 Moles n Defined as the number of carbon atoms in exactly 12 grams of carbon- 12. n 1 mole is 6.02 x particles. n Treat it like a very large dozen n 6.02 x is called Avogadro's number.

4 4 Representative particles n The smallest pieces of a substance. n For an element it is an atom. –Unless it is diatomic n For a molecular compound it is a molecule. n For an ionic compound it is a formula unit.

5 5 Conversion factors n Used to change units. n Three questions –What unit do you want to get rid of? –Where does it go to cancel out? –What can you change it into?

6 6 Calculation question n How many molecules of CO 2 are the in 4.56 moles of CO 2 ?

7 7 Calculation question n How many moles of water is 5.87 x molecules?

8 8 Calculation question n How many atoms of carbon are there in 1.23 moles of C 6 H 12 O 6 ?

9 9 Measuring Moles n The amu was one twelfth the mass of a carbon 12 atom. n Since the mole is the number of atoms in 12 grams of carbon-12, n the decimal number on the periodic table is –The mass of the average atom in amu –the mass of 1 mole of those atoms in grams.

10 10 Gram Atomic Mass n The mass of 1 mole of an element in grams. n grams of carbon has the same number of atoms as 1.01 grams of hydrogen and grams of iron. n We can write this as g C = 1 mole n We can count things by weighing them.

11 11 Examples n How much would 2.34 moles of carbon weigh?

12 12 Examples n How many moles of magnesium in 4.61 g of Mg?

13 13 Examples n How many atoms of lithium in 1.00 g of Li?

14 14 Examples n How much would 3.45 x atoms of U weigh?

15 15 What about compounds? n in 1 mole of H 2 O molecules there are two moles of H atoms and 1 mole of O atoms n To find the mass of one mole of a compound –determine the moles of the elements they have –Find out how much they would weigh –add them up

16 16 What about compounds? n What is the mass of one mole of CH 4 ? n 1 mole of C = g n 4 mole of H x 1.01 g = 4.04g n 1 mole CH 4 = = 16.05g

17 17 Molar Mass n The mass of 1 mole n What is the molar mass of Fe 2 O 3 ? n 2 moles of Fe x g = g n 3 moles of O x g = g n The GFM = g g = g

18 18 Calculate the molar mass of the following n C 6 H 12 O 6 n (NH 4 ) 3 PO 4

19 19 Using Molar Mass Finding moles of compounds Counting pieces by weighing

20 20 Molar Mass n The number of grams in 1 mole of atoms, formula units, or molecules. n We can make conversion factors from these. n To change grams of a compound to moles of a compound. n Or moles to grams

21 21 For example n How many moles is 5.69 g of NaOH? l need to change grams to moles l for NaOH l 1mole Na = 22.99g 1 mol O = g 1 mole of H = 1.01 g l 1 mole NaOH = g

22 22 For example n How many moles is 5.69 g of NaOH? l need to change grams to moles l for NaOH l 1mole Na = 22.99g 1 mol O = g 1 mole of H = 1.01 g l 1 mole NaOH = g

23 23 Gases and the Mole

24 24 Gases n Many of the chemicals we deal with are gases. n They are difficult to weigh, so well measure volume n Need to know how many moles of gas we have. n Two things affect the volume of a gas n Temperature and pressure n Compare at the same temp. and pressure.

25 25 Standard Temperature and Pressure n Avogadro's Hypothesis - at the same temperature and pressure equal volumes of gas have the same number of particles. n 0ºC and 1 atmosphere pressure n Abbreviated atm n 273 K and kPa n kPa is kiloPascal

26 26 At Standard Temperature and Pressure n abbreviated STP n At STP 1 mole of gas occupies 22.4 L n Called the molar volume n Used for conversion factors n Moles to Liter and L to mol

27 27 Examples n What is the volume of 4.59 mole of CO 2 gas at STP?

28 28 Density of a gas n D = m /V n for a gas the units will be g / L n We can determine the density of any gas at STP if we know its formula. n To find the density we need the mass and the volume. n If you assume you have 1 mole than the mass is the molar mass (PT) n At STP the volume is 22.4 L.

29 29 Examples n Find the density of CO 2 at STP.

30 30 Quizdom n Find the density of CH 4 at STP.

31 31 The other way n Given the density, we can find the molar mass of the gas. n Again, pretend you have a mole at STP, so V = 22.4 L. n m = D x V n m is the mass of 1 mole, since you have 22.4 L of the stuff. n What is the molar mass of a gas with a density of g/L?

32 32 All the things we can change

33 33 Volume Ions Atoms Representative Particles Mass PT Moles 6.02 x L Count

34 34 Percent Composition n Like all percents n Part x 100 % whole n Find the mass of each component, n divide by the total mass.

35 35 Example n Calculate the percent composition of a compound that is 29.0 g of Ag with 4.30 g of S.

36 36 Getting it from the formula n If we know the formula, assume you have 1 mole. n Then you know the pieces and the whole.

37 37 Examples n Calculate the percent composition of C 2 H 4 ?

38 38 Examples n What is the percent composition of Aluminum carbonate.

39 39 Percent to Mass n Multiply % by the total mass to find the mass of that component. n How much aluminum in 450 g of aluminum carbonate?

40 40 Empirical Formula From percentage to formula

41 41 The Empirical Formula n The lowest whole number ratio of elements in a compound. n The molecular formula the actual ratio of elements in a compound. n The two can be the same. n CH 2 empirical formula n C 2 H 4 molecular formula n C 3 H 6 molecular formula n H 2 O both

42 42 Finding Empirical Formulas n Just find the lowest whole number ratio n C 6 H 12 O 6 n CH 4 N 2 n It is not just the ratio of atoms, it is also the ratio of moles of atoms.

43 43 Calculating Empirical Formulas n Means we can get ratio from percent composition. n Assume you have a 100 g. n The percentages become grams. n Turn grams to moles. n Find lowest whole number ratio by dividing everything by the smallest moles.

44 44 Example n Calculate the empirical formula of a compound composed of % C, % H, and %N. n Assume 100 g so n g C x 1mol C = mole C gC n g H x 1mol H = 16.1 mole H 1.01 gH n g N x 1mol N = mole N gN

45 45 Example n The ratio is mol C = 1 mol C molN 1 mol N n The ratio is 16.1 mol H = 5 mol H molN 1 mol N nC1H5N1nC1H5N1nC1H5N1nC1H5N1 n Caffeine is 49.48% C, 5.15% H, 28.87% N and 16.49% O. What is its empirical formula?

46 46 Empirical to molecular n Caffeine is 49.48% C, 5.15% H, 28.87% N and 16.49% O. What is its empirical formula? n Since the empirical formula is the lowest ratio the actual molecule would weigh the same or more. n By a whole number multiple. n Divide the actual molar mass by the the mass of one mole of the empirical formula. n You will get a whole number. n Multiply the empirical formula by this.

47 47 Example n n A compound has an empirical formula of ClCH 2 and a molar mass of g/mol. What is its molecular formula? n n A compound has an empirical formula of CH 2 O and a molar mass of g/mol. What is its molecular formula?

48 48 Percent to molecular n Take the percent x the molar mass –This gives you mass in one mole of the compound n Change this to moles –You will get whole numbers –These are the subscripts n Caffeine is 49.48% C, 5.15% H, 28.87% N and 16.49% O. It has a molar mass of 194 g. What is its molecular formula?

49 49 Example n n Ibuprofen is % C, 8.80 % H, % O, and has a molar mass of about 207 g/mol. What is its molecular formula?


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