Warm Up. What is a mole? What is molar mass? What is Avogadro s number?

Transcription

1 Warm Up What is a mole? What is molar mass? What is Avogadro s number?

2 Chapter 7 The Mole and Chemical Compostion

3 Unit Essential Question: HOW CAN CHEMICAL COMPOSITION BE DETERMINED?

4 Lesson Essential Question: HOW IS THE MOLE USED IN CONVERSIONS?

5 Section 1: Avogadro s Number and Molar Conversions 1 mole = x particles SI unit for amount of substance. It s a counting unit (like a dozen). Remember that the unit of particles can be: ions, molecules (mcs.), atoms, formula units (f.u.), etc. covalent compounds ionic compounds Recall that formula units = simplest ratio of ions in an ionic compound.

6 Recall your mole map!

7 Converting moles ó particles Same as Chapter 3, but it will involve molecules, formula units, or ions instead of just atoms. Steps: 1) Need 1mol = x10 23 molecules, etc. 2) Use dimensional analysis- turn this into a fraction! *Be sure to place the correct units on the top and bottom so they cancel!

8 Sample Problems 1 & 2: Moles & Particles Find the number of molecules in 2.5 mol of sulfur dioxide. 1.5 x molecules SO 2 A sample contains 3.01 x molecules of sulfur dioxide. Determine the amount in moles mol SO 2

9 Molar Mass Amount of mass (in grams) in 1 mole of a substance. Use molar masses from the periodic table. Round to 2 decimal places! Use units of g/mol. Example: C: 12.01g/mol means that 1 mol C = g Cl: 35.45g/mol means that 1 mol Cl = 35.45g Use to convert between moles and mass.

10 Sample Problems 3 & 4: Moles & Mass What is the mass of 5.3mol Be? 48g Be If you have 27.0g of manganese, how many moles do you have? 0.491mol Mn

11 Molar Masses of Compounds Add together the molar masses of all elements or ions present. Ex: CH 4 C: 12.01g/mol H: 1.01g/mol 12.01g/mol + 4(1.01g/mol) = 16.05g/mol This means that 1 mole of CH 4 has a mass of 16.05g. You will need to calculate the molar mass of a compound whenever you are converting between mass and moles!

12 Additional Molar Mass Examples: Element Ag = g/mol Diatomic Element/molecule Br 2 = x 2 = g/mol Molecule (Covalent compound) H 2 O = (1.01 x 2) = g/mol Formula unit (Ionic compound) Ca(NO 3 ) 2 = (2 x 14.01) + (6 x 16.00) = g/mol

13 Sample Problem 5: Mass to Moles with a Compound Find the number of moles present in 47.5 g of glycerol, C 3 H 8 O 3. Hint: you will need to calculate the molar mass of glycerol! Glycerol s molar mass: 92.11g/mol 0.516mol C 3 H 8 O 3

14 Sample Problem 6: Number of Particles to Mass Remember- you can t go directly between mass (g) and the number of particles! You must convert to moles first! Find the mass in grams of 2.44 x atoms of carbon g C

15 More Practice How many moles of iron (III) sulfate, Fe 2 (SO 4 ) 3, are there in a 178g sample? 0.445mol

16 Lesson Essential Question: HOW ARE MOLAR MASSES ON THE PERIODIC TABLE DETERMINED?

17 Example #3: N 2 O 3 For every 1mol of N 2 O 3 there are 2mol of N atoms and 3mol of O atoms. Mole Ratios in Chemical Formulas Ratios can be formed between amounts of elements or ions within a compound. Look at the subscripts. Example #1: CaCl 2 For every 1mol of CaCl 2 there is 1mol of Ca +2 ions and 2mol of Cl - ions. Example #2: Na 2 CO 3 For every 1mol of Na 2 CO 3, there are 2mol of Na + ions and 1mol of CO 3-2 ions.

18 Practice If you have one mole of strontium cyanide, Sr(CN) 2, how many moles of strontium ions are there? How many moles of cyanide ions are there? Given the compound P 2 O 5 what is the mole ratio of P atoms to O atoms?

19 Section 2: Relative Atomic Mass and Chemical Formulas Periodic table masses are averages of all isotopes present. Recall that we said a weighted average is used- takes into account the amount of each isotope. Average atomic mass: (% x atomic mass)+(% x atomic mass)+ 100 Note: % is the percent abundance (how often the element is found as that isotope in nature).

20 Sample Problem The mass of a Cu-63 atom is amu, and that of a Cu-65 atom is amu. If the abundance of Cu-63 is 69.17% and the abundance of Cu-65 is 30.83%, what is the average atomic mass of copper?

21 Lesson Essential Questions: WHAT INFORMATION CAN BE DETERMINED FROM FORMULAS? HOW CAN FORMULAS BE DETERMINED?

22 Calculating Percent Composition Tells you the percent each element makes up of the whole compound. Step 1: Determine the molar mass of the entire compound. Step 2: Divide each element s total molar mass by the molar mass of the compound. Step 3: Multiply by 100 to get percent. Step 4: Check your answer by adding up the percentages to makes sure they equal 100%.

23 Percent Composition Cont. Calculating the percent composition of a compound can be helpful in determining the formula/identity. Example: Iron and oxygen form two compounds: Fe 2 O 3 and FeO Fe 2 O 3 = 69.9% Fe and 30.1% O FeO = 77.7% Fe and 22.3% O

24 Sample Problem #I Calculate the percent composition of copper (I) sulfide. Sample Problem #2 Calculate the percent composition of isopropyl alcohol, (CH 3 ) 2 CHOH.

25 Determining Empirical Formulas The empirical formula shows the simplest ratio of elements/ions in the compound. Ionic compounds are represented with empirical formulas. Given percent composition data, you can determine the empirical formula of a compound. Step 1: Assume 100 g of the sample- put g in for %. Ex: 18.2% O à 18.2g Step 2: Convert grams to moles. Step 3: Divide each mole value by the smallest mole value. This will tell you the number of each element that appears in the formula.

26 Determining Empirical Formulas Cont. Step 4: If you get a decimal, multiply ALL numbers by a whole number to turn the decimal into a whole number. The numbers you will need to multiply by should be relatively small (2, 3, etc.)

27 Sample Problem #1 Chemical analysis of a liquid shows that it is 60.0% C, 13.4% H, and 26.6% O by mass. Calculate the empirical formula of this substance. Sample Problem #2 A compound is found to contain 38.77% Cl and 61.23% O. What is the empirical formula?

28 Molecular Formulas Show the actual numbers of elements in the compound- not necessarily the simplest formula. Often seen for covalent compounds. They will be a whole number multiple of the empirical formula (can t be a decimal). In other words: n(empirical formula) = molecular formula where n is a whole number. Ex: 6(CH 2 O) à C 6 H 12 O 6 Molecular and empirical formulas can be the same!

29 Molecular Formulas Cont.

30 Molecular Formulas Cont. The molecular formula can be determined from the empirical formula and experimental molar mass of a compound. Step 1: Determine the molar mass of the given empirical formula. Step 2: Solve for n by dividing the experimental molar mass by the molar mass of the empirical formula. *Remember: n(empirical formula) = molecular formula Step 3: Multiply the subscripts in the empirical formula by n.

31 Sample Problem #1 The empirical formula for a compound is P 2 O 5. Its experimental molar mass is 284g/mol. Determine the molecular formula of the compound. Sample Problem #2 A brown gas has the empirical formula NO 2. Its experimental molar mass is 46g/ mol. What is the molecular formula?

32 Hydrates- Honors Only Not in the textbook. Hydrates ionic compounds that contain water molecules within the crystal structure. Example: CuSO 4 5H 2 O Anhydrous without the water = CuSO 4

33 Determining Hydrate Formulas Formula can be determined if given: the mass of the hydrate, the anhydrous mass, and the formula of the ionic compound. Step 1: Determine the mass of water in the hydrate (subtract anhydrous mass). Step 2: Convert the anhydrous ionic compound mass and water mass to moles. Step 3: Divide both molar amounts by the smallest number. This gives you the number of water molecules in the hydrate.

34 Sample Problem #1 A 5.82 g sample of Mg(NO 3 ) 2 XH 2 O in an evaporating dish is heated until it is dry. The mass of the anhydrous sample is 2.63 g Mg(NO 3 ) 2. What is the formula for the hydrate?

35 Determining % Water in a Hydrate Formula can be determined if given the formula of the hydrate. Step 1: Calculate the mass of the entire hydrate and the mass of just the water. Step 2: Divide the mass of the water by the mass of the entire hydrate and multiply by 100 to get a percent.

36 Sample Problem #2 What percentage, by mass, of water is found in the hydrate CuSO 4 5H 2 O?