Chapter 8 Chemical Quantities

Transcription

1 Chapter 8 Chemical Quantities Introductory Info The atomic masses of the elements on the periodic table are in the units. These measurements are based on the mass of the standard isotope of the element, which has a mass of. All other elements have a mass that is a fraction or multiple of that isotope. Hydrogen has a mass that is 1/12 of the standard isotope so hydrogen has an atomic mass of. Magnesium s mass is approximately twice the standard isotope, giving it a mass of. One atomic mass unit is extremely small being equal to the mass of one atom ( (73) kg), so it is impractical to use this value in normal lab calculations, so another, larger unit is needed. It should be noted that the ratio of masses of atoms is constant: the ratio of one atom of carbon to one atom of magnesium is to _. The ratio of 2 atoms of carbon to 2 atoms of magnesium is _ to _. The ratio of 10 atoms of carbon to 10 atoms of magnesium is _ to _. The ratio of and number of carbon atoms to the same number of magnesium atoms is to _. The common practice in laboratory calculations is to replace the amu value for an element with grams, giving us a gram atomic mass. It has been found experimentally that the gram atomic mass of an element contains atoms. This is known as Avogadro s number and is known as a mole. A mole of an element is the mass in grams of atoms of that element. One atom of carbon has a mass of amu. Therefore 6.02 x atoms of carbon have a mass of grams. One atom of magnesium has a mass of amu. Therefore 6.02 x atoms of magnesium have a mass of grams. When elements combine to form compounds, the particles of that compound are known as if the compound is covalently bonded, or, if the compound is ionically bonded. The mass of one mole of a compound is the gram formula mass (also known as the molar mass) and is the sum of the gram atomic masses of the atoms in the compound. Thus, a mole of water (H 2 O) consists of _ molecules and has a mass of grams 1 mole HCl = molecules = _grams 2 moles H 2 = molecules = atom = _ grams 1 molecule H 2 O 2 = total atoms 1 mole H 2 O 2 = _ atoms H and _ atoms O 1 mole FeCl 3 = moles Fe +3 and moles Cl - 1 mole FeCl 3 = grams Fe +3 and grams Cl - 1 mole FeCl 3 = total ions and moles of ions H. Cannon, C. Clapper and T. Guillot Klein High School

2 Molecular Weight and Moles Find the molar mass or formula mass of each of the following 1. HNO 3 2. Ammonium nitrate 3. Fe 2 O 3 4. Rubidium Sulfite 5. H 3 PO 4 6. Lithium Carbonate 7. K 2 SO 4 8. Magnesium Hydroxide 9. Be 5 As Aluminum Sulfate Find the mass of each of the following expressed in grams mol of HC 2 H 3 O mol of sodium acetate mol of K 2 CrO g of calcium phosphate mol of Ca(ClO 3 ) mol lithium sulfate mol of Ba(NO 3 ) mol of iron (II) nitrate mol of Na 2 Cr 2 O mol copper (II) acetate Find the mass of each of the following, expressed in moles g of CaCO g of arsenic trichloride g of Ni(NO 3 ) g of calcium phosphide g of C 6 H 12 O g of calcium acetate g K 3 PO g of aluminum nitrate g of Bi(OH) g of iron(ii)phosphate Calculate grams for each of the following x atoms of Na x ions on NaOH x10 23 atoms of Ca x ions Na x atoms of S x molecules H 2 O Calculate the mass in grams for each of the following moles Na moles of O moles Ca moles of Al moles Mg moles H 2 8-2

3 moles Cl moles H 2 SO moles CaCO moles KI moles MgCl moles Ca(OH) moles Al 2 O moles Ca(NO 2 ) 2 Calculate the number of moles of each of the following g F g of Zn g Li g Br g Ne g Fe g Ca g SO g NaOH g Na 2 S g MgCO g K 2 SO g ZnO g H 2 O 2 Calculate the number of atoms, molecules or ions for each of the following moles Na atoms g S atoms mole N atoms g Ca atoms g Na atoms moles CO 2 molecules moles K + ions g H 2 O molecules g H 2 S molecules moles Mg +2 ions 8-3

4 Chemical Quantities Exercise #2 Mole/ Gram Problems 1. How many grams are in 7.20 moles of dinitrogen trioxide? 2. Find the number of moles in 922 g of iron(iii)oxide. 3. Calculate the number of grams in 2.4 moles of potassium oxalate. 4. Calculate the number of moles in 450 g of barium silicate. 5. How many moles are in 1206 g of calcium sulfate? g of chromium (III) bromide contains how many moles? 7. Calculate the number of moles in g of ammonium acetate. 8. How many grams are in 4.5 mol of hydrogen phosphate? mol of iron(ii)permanganate has what mass? 10. What mass of sodium carbonate would be equal to 1.24 mol? 11. What is the mass of mol of potassium nitrate? 12. What is the formula mass of strontium chloride? 13. What is the formula mass of sodium carbonate? 14. How many grams of sodium sulfate are in 5.4 moles of sodium sulfate? 15. How many moles are in 560 g of calcium phosphate? 16. How many moles are in 5.6 x 10 3 g of ammonium oxalate? 17. How many g of iron (III)perchlorate are in 625 moles? g of plumbous silicate was used in an experiment. Calculate the number of moles contained in the amount of this substance. 19. If 4.57 moles of ammonium acetate were used to activate a chemical reaction, how many grams were needed? 20. Determine which substance provides the most mass; 3.5 moles of iron(iii) carbide or 3.5 moles of iron (III) oxide 21. Calculate the number of moles contained in 5.6 x 10 5 grams of silver nitrate 22. Determine the number of moles in 450g of potassium dichromate. 23. Which has more mass, 3.5 moles of calcium carbonate or 3.5 moles of calcium phosphate? 24. Which has more moles, 100.0g of sodium hydroxide or g of potassium hydroxide? 25. What mass of water would it take to give you one mole? 8-4

5 Practice Problems For the following determine moles or grams 1. moles in 56g of calcium phosphate 2. grams in 5.6 moles of silver nitrate 3. grams in 6.7 moles of zinc acetate 4. moles in g of ferric sulfide 5. moles in 1.2 x 10 4 g of mercurous nitrate 6. moles in g of potassium oxalate 7. grams in moles of copper (II) sulfate 8. grams in 1.2 x 10-2 moles aluminum silicate 9. grams in moles of mercury (I) bicarbonate 10. moles in 5.6 x 10 4 g of magnesium chlorate Chemical Quantities Exercise #4 For the following calculate the gram formula mass 11. aluminum acetate 12. calcium nitrate 13. nitrogen trioxide 14. mercuric chloride 15. silicon tetraflouride 16. tin (IV) sulfite 17. ammonium hydroxide 18. strontium acetate 19. Calcium hydroxide 20. Hydrogen phosphate For the following calculate number of particles moles of sodium chloride grams of barium chloride moles of silver sulfate grams of aluminum bromide moles of calcium flouride 8-5

6 Exercise #5 Percent Composition 1. What is the percent composition of H in H 2 O? 2. What is the % composition of Al in AlPO 4? 3. What is the % composition of Ca in Ca(OH) 2? 4. What is the % composition of C in Ba(C 2 H 3 O 2 ) 2? 5. What is the % of Pb in PbCl 4? 6. What is the % composition of NH 4 in (NH 4 ) 2 SO 3? 7. What is the % composition of hydrogen in calcium hydroxide? 8. What is the % composition of oxygen in CO 2 9. A 5.2 g piece of magnesium combines with 3.1g of oxygen to form a compound, what is the % composition of magnesium in this compound? 10. A 2.2g piece of sodium combines with 6.3g of iodine to form a compound. What is the % composition of iodine in this compound? 11. A 15.2g piece of sulfur combines with 9.4g of potassium to form a compound. What is the % of potassium in the compound? 12. A 1.5g piece of barium combines with 9.0g of phosphorous to form a compound. What is the % composition of barium in this compound? 13. An 8.20g piece of magnesium combines completely with 5.40g of chlorine to form a compound. What is the % composition of this compound? 14. Calculate the % composition of carbon in calcium acetate. 15. Calculate the percent of sulfur in sodium bisulfate. 16. Which compound contains the most hydrogen; a. 20.0g of potassium hydrogen sulfate b. 124g of calcium acetate c. 378g of hydrogen cyanide 17. If 500g of sodium chloride were analyzed, how many grams of sodium metal would be found? 18. How many grams of hydrogen can be derived from 5.6 x 10 5 g of water? 19. Iron metal can be extracted from iron ore. There are two types of iron ore, iron (III) oxide and iron (II) oxide. If you had 800.0g of each ore, which ore would produce the most iron metal? 20. How many moles in 30.0 cm 3 of copper? (the density of copper is 8.92g/cm 3 ) 8-6

7 Empirical Formula _ 1. What is the empirical formula of a compound containing 63g of Rb and 5.9g of O? _ 2. What is the empirical formula of a compound with 0.159g of U and 0.119g of Cl? _ 3. Write the empirical formula for the compound that has 7.22g of Ni, 2.53g of P, and 5.25g of O. _ 4. If a compound contains 0.285g of Ca, 0.236g of S and 0.469g of O, what is its empirical formula? _ 5. If a compound is made up of 32.8% Cr and 67.2% Cl, what is it s empirical formula? _ 6. What is the empirical formula of a compound found to contain 42.7% Co and 57.3% Se? _ 7. Find the empirical formula of a compound that is 56.6% La and 43.4% Cl. _ 8. What is the empirical formula of a compound that is Ta and 18.1% O? _ 9. If the % composition of a compound is 92.3% C and 7.7% H, what is the empirical formula? _ 10. Find the empirical formula of a compound with a % composition of 26.7% P, 12.2% N and 61.2% Cl. _ 11. A hydrate was analyzed to determine that it contained 5.262g of Tl(NO 3 ) 3 and 0.789g of water, what is the formula of the hydrate? _ 12. What is the formula for a hydrate if it contains 2.94g of Sn(NO 3 ) 2 and 4.37g of water? _ 13. In a chemical reaction, 1.58g of copper combine with sulfur to give 1.98g of a new compound. What is the empirical formula of this new compound? _ 14. An analysis determines that a compound is made of 42.9% carbon and 57.1% oxygen, what is the empirical formula? _ 15. In a chemical process the following data was collected. 44.5% copper and 55.7% bromine. Calculate the correct empirical formula for the compound. _ 16. From the following data determine the correct empirical formula. _ 38.9% Ba, 29.4% Cr and 31.7% O. 17. In a chemical reaction, 0.274g of aluminum combine with iodine to form 4.41g of a new product. What is the empirical formula of this product? _ 18. What is the empirical formula of a compound containing 9.93% carbon, 58.6% chlorine and 31.1% fluorine? 8-7

8 Molecular Formula 1. The molecular formula of a compound is either the same as its empirical formula or a _ of it. 2. What do you need to know to calculate the molecular formula of a compound? 3. If you divide the molecular mass of a compound by the empirical formula mass, what is the result? 4. What would you use to convert the empirical formula of a compound to a molecular formula? 5. Gas X is found to be 24.0% carbon and 76.0% fluorine, what is its empirical formula? Given that the molar mass of gas X is g/mol, determine its molecular formula. 6. Ribose is an important sugar (part of RNA), with a molar mass of g/mol. If its empirical formula is CH 2 O, what is its molecular formula? 7. Naphthalene is a soft covalent solid that is often used in mothballs. Its molar mass is g/mol and it contains 93.75% carbon and 6.25% hydrogen. Determine the molecular formula of naphthalene from this data. 8. What is the molecular formula for a compound with a molecular mass of 30.0 g/mol containing 80.0% hydrogen and 20.0 % oxygen? 9. If a compound contains 37.8% carbon 6.3% hydrogen, 55.8% chlorine and a molecular mass of 127.0g/mol, what is the molecular formula 8-8

9 Chemical Quantities Review 1. What is the molecular mass of dinitrogen pentoxide? 2. Calculate the formula weight of aluminum silicate. 3. What is the % composition of phosphorous in plumbic phosphate? 4. What is the % composition of each element in barium acetate? 5. What is the mass of 3.04 moles of sodium carbonate? 6. How many atoms are there in 84.4g of carbon 7. How many moles are in 481g of magnesium phosphate? 8. What is the mass of 5.94 x formula units of calcium chlorate? 9. A 5.2g piece of magnesium combines with 3.1g of oxygen, what is the % composition of magnesium in the compound? 10. If 15.2g of sulfur combine with 9.4g of potassium, what is the % composition of potassium in the compound? 11. What is the % composition of a compound formed when 6.3g of iodine combine w 2.2g of sodium? 12. What is the empirical formula for a compound that contains 23.6g of sulfur, 46.9g oxygen and 29.5g of calcium? 13. Find the empirical formula for the compound with the following analysis: 8.54g M 16.78g C, 2.1g H, 22.49g O. 14. Find the empirical formula for the compound with the following analysis: 9.93% carbon, 58.6% chlorine, and 31.3% fluorine. 15. What is the empirical formula for a compound that contains 5.03g of aluminum and 8.97g of sulfur? 16. What is the percent composition of magnesium in 3.50 mole of magnesium nitride? 17. A compound is found to contain 33.3% calcium, 40.0% oxygen, and 26.7% sulfur. What is the empirical formula for the compound? 18. What is the molecular formula of a compound with a molecular mass of 78 g/mol if the % composition is 92.3% carbon, and 7.7% hydrogen? 19. A compound is found to have a percent composition of 80.4% bismuth, 18.5% oxygen and 1.16% hydrogen. What is the empirical formula? 8-9

10 Lab #3 Chapter 8 LAB: DETERMINATION OF THE FORMULA OF A HYDRATE OBJECTIVE: Given a sample of hydrated salt and appropriate apparatus, the student will be able to experimentally determine the correct formula for the compound. APPARATUS: goggles, Bunsen burner, matches, crucible, triangle, crucible tongs, ringstand and ring, spatula, balance. PROCEDURE: 1. Weigh a crucible to the nearest 0.01 gram. Add approximately 2 grams of salt. Record total weight of salt and crucible to the nearest 0.01 gram. Subtract to find the weight of your salt. 2. Support the crucible on the triangle at an appropriate height to receive the maximum heat of the flame. WARNING: hydrates, upon heating, can decrepitate (explode). Protect your eyes at all times. Heat the crucible gently at first. Too rapid heating may cause loss by spattering because the water of hydration is driven off too rapidly. Gradually heat more strongly, keeping the crucible in the hottest part of the flame. After ten minutes of strong heating, gradually withdraw the flame, wait a moment, then use the crucible tongs to move the crucible to the base of the ringstand to cool for at least three minutes. Never put a hot crucible on a balance. When cooled enough to touch, move to the balance with crucible tongs, weigh and record. 3. Repeat the heating, cool, and reweigh. Continue this process until you get two successive weighing within 0.20 grams. This is called heating to constant weight. Discard the used, now anhydrous salt in the trash. The crucible may then be washed out with tap water. 4. SHOW ALL CALCULATIONS CLEARLY AND NEATLY. WATCH SIGNIFICANT DIGITS. a. Use the lowest weight to find the mass of anhydrous salt. b. Find the formula mass of your salt. c. Find moles of your salt. d. Subtract the two masses to find the mass of water e. Find moles of water f. Right the formula for the hydrate g. Calculate the percentage of water in the hydrated salt h. Find the formula mass of the hydrated salt. i. Calculate the theoretical percentage of water in the hydrate j. Calculate your percent error. 8-10