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Chemical Mathematics Atomic Mass Atoms are so small, it is difficult to discuss how much they weigh in grams. Use atomic mass units. An atomic mass unit.

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Presentation on theme: "Chemical Mathematics Atomic Mass Atoms are so small, it is difficult to discuss how much they weigh in grams. Use atomic mass units. An atomic mass unit."— Presentation transcript:

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2 Chemical Mathematics

3 Atomic Mass Atoms are so small, it is difficult to discuss how much they weigh in grams. Use atomic mass units. An atomic mass unit (amu) is one twelth the mass of a carbon-12 atom. This gives us a basis for comparison. The decimal numbers on the table are atomic masses in amu.

4 They are not whole numbers Because they are based on averages of atoms and of isotopes. Can figure out the average atomic mass from the mass of the isotopes and their relative abundance. Add up the percent as decimals times the masses of the isotopes.

5 Examples There are two isotopes of carbon 12 C with a mass of amu(98.892%), and 13 C with a mass of amu (1.108%). There are two isotopes of nitrogen, one with an atomic mass of amu and one with a mass of amu. What is the percent abundance of each? All of these are solved with algebra !

6 The Mole The mole is a number. A very large number, but still, just a number x of anything is a mole Similar in idea to a dozen. – A dozen is always 12 items – A mole is always 6.02 x items A mole is in exactly 12 grams of carbon-12. Makes the numbers on the table the mass of the average atom.

7 The Mole and the Magic Formula From a practical standpoint, we want to be able to convert grams to moles and moles to grams. THIS IS OF UTMOST IMPORTANCE IN CHEMISTRY !! We need to introduce the most used formula in Chemistry. N = g / M Moles = grams / formula wt.

8 Molar mass Mass of 1 mole of a substance. Often called molecular weight. To determine the molar mass of an element, look on the table. To determine the molar mass of a compound, add up the molar masses of the elements that make it up.

9 Some Generic Names You can use the terms weight and mass interchangeably here. The most general term is FORMULA WT. Molar mass is another common term. Atomic weight is the weight of an atom. Molecular weight is the weight of a molecule. Ionic weight is the weight of an ion Etc., etc., etc.

10 Find the molar mass of CH 4 Mg 3 P 2 Ca(NO 3 ) 3 Al 2 (Cr 2 O 7 ) 3 CaSO 4 · 2H 2 O

11 Percent Composition Percent of each element a compound is composed of. Find the mass of each element, divide by the total mass, multiply by a 100. If a percent is given instead, assume a 100 gram sample and simply change the percents to grams !! Find the percent composition of CH 4 Al 2 (Cr 2 O 7 ) 3 CaSO 4 · 2H 2 O

12 Working backwards From percent composition, you can determine the empirical formula. Empirical Formula the lowest ratio of atoms in a molecule. Based on mole ratios. A sample is 59.53% C, 5.38% H, 10.68% N, and 24.40% O What is its empirical formula ?

13 Empirical Formula A gram sample of a compound (vitamin C) composed of only C, H, and O is burned completely with excess O g of CO 2 and g of H 2 O are produced. What is the empirical formula? Follow the steps !

14 Empirical To Molecular Formulas Empirical is lowest ratio. Molecular is actual molecule. Need Formula weight. Ratio of empirical to molar mass will tell you the molecular formula. Must be a whole number because molecules are made up of whole number of atoms.

15 How to determine Empirical Formulas No matter what unit is given, you need it to be in moles. If moles, you are ready. If grams, use n = g/M If a %, assume 100 grams and use n = g/M After moles are obtained, you have a proper ratio. Simply make sure that it is in whole numbers !

16 Example A compound is made of only sulfur and oxygen. It is 69.6% S by mass. Its molar mass is 184 g/mol. What is its formula?

17 Chemical Equations Are chemical sentences. Describe what happens in a chemical reaction. Reactants  Products Equations should be balanced. Have the same number of each kind of atoms on both sides because...

18 Balancing equations CH 4 + O 2  CO 2 + H 2 O ReactantsProducts C11 O23 H42

19 Balancing equations CH 4 + O 2  CO H 2 O ReactantsProducts C11 O23 H424

20 Balancing equations CH 4 + O 2  CO H 2 O ReactantsProducts C11 O23 H424 4

21 Balancing equations CH 4 + 2O 2  CO H 2 O ReactantsProducts C11 O23 H

22 Abbreviations (s)  (g)  (aq) heat  catalyst

23 Practice Ca(OH) 2 + H 3 PO 4  H 2 O + Ca 3 (PO 4 ) 2 Cr + S 8  Cr 2 S 3 KClO 3 (s)  Cl 2 (g) + O 2 (g) Solid iron(III) sulfide reacts with gaseous hydrogen chloride to form solid iron(III) chloride and hydrogen sulfide gas. Fe 2 O 3 (s) + Al(s)  Fe(s) + Al 2 O 3 (s)

24 Meaning A balanced equation can be used to describe a reaction in molecules and atoms. NEVER GRAMS ! Chemical reactions happen molecules at a time or dozens of molecules at a time or moles of molecules.

25 Stoichiometry Given an amount of either starting material or product, determining the other quantities. Use conversion factors from – molar mass (g - mole) – balanced equation (mole - mole) Keep track by setting up proportions.

26 Examples One way of producing O 2 (g) involves the decomposition of potassium chlorate into potassium chloride and oxygen gas. A 25.5 g sample of Potassium chlorate is decomposed. How many moles of O 2 (g) are produced ? How many grams of potassium chloride ? How many grams of oxygen ?

27 Examples A piece of aluminum 5.11 cm. by 3.23 cm. x cm. is dissolved in excess HCl(aq). How many grams of H 2 (g) are produced? How many grams of each reactant are needed to produce 15 grams of iron form the following reaction? Fe 2 O 3 (s) + Al(s)  Fe(s) + Al 2 O 3 (s)

28 Examples K 2 PtCl 4 (aq) + NH 3 (aq)  Pt(NH 3 ) 2 Cl 2 (s)+ KCl(aq) what mass of Pt(NH 3 ) 2 Cl 2 can be produced from 65 g of K 2 PtCl 4 ? How much KCl will be produced? How much from 65 grams of NH 3 ?

29 Stoichiometric Yield How much you get from an chemical reaction

30 Limiting Reagent Reactant that determines the amount of product formed. It governs the reaction. The one you run out of first. Makes the least product. If it is not easily apparent which reactant is the limiting reagent, you can always set up 2 proportions and see which one fits. One will fit and the other one won’t – ALWAYS.

31 Limiting reagent To determine the limiting reagent requires that you do two stoichiometry problems. Figure out how much product each reactant makes. The one that makes the least is the limiting reagent.

32 Example Ammonia is produced by the following reaction N 2 + H 2  NH 3 What mass of ammonia can be produced from a mixture of 100. g N 2 and 500. g H 2 ? How much unreacted material remains?

33 Excess Reagent The reactant you don’t run out of. The amount of stuff you make is the yield. The theoretical yield is the amount you would make if everything went perfect. The actual yield is what you make in the lab.

34 Percent Yield % yield = Actual x 100% Theoretical Or stated another way % yield = what you got x 100% what you should have got

35 Examples Aluminum burns in bromine producing aluminum bromide. In a laboratory 6.0 g of aluminum reacts with excess bromine g of aluminum bromide are produced. What is the percent yield.

36 Examples Years of experience have proven that the percent yield for the following reaction is 74.3% Hg + Br 2  HgBr 2 If 10.0 g of Hg and 9.00 g of Br 2 are reacted, how much HgBr 2 will be produced? If the reaction did go to completion, how much excess reagent would be left?

37 Examples Commercial brass is an alloy of Cu and Zn. It reacts with HCl by the following reaction Zn(s) + 2 HCl(aq)  ZnCl 2 (aq) + H 2 (g) Cu does not react. When g of brass is reacted with excess HCl, g of ZnCl 2 are eventually isolated. What is the composition of the brass ? First think of what you are solving for. Plan the problem out.


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