Percentage composistion 2024. Power point

Percentage composition

1. QUANTITATIVE CHEMISTRY Percent composition and chemical formulae

2. PDN 1. What is a chemical formula? 2. Write the types of chemical formula do you know and give their names. 3. How do you determine a molecular formula? 2

3. By the end of the lesson learners should be able to Calculate percentage composition of a substance Define empirical formula and molecular formula. Differentiate an empirical formula from a molecular formula. Determine an empirical formula from given percentages or mass. Determine a molecular formula 3 Lesson objectives

4. Percent Composition of Compounds The percent composition of a compound is the percent by mass of each element in the compound. The composition of compounds is given by their formulas. For example, the formula of carbon dioxide ( ), tells us it is composed of carbon and oxygen.   Every molecule of is composed of one atom of carbon and two atoms of oxygen.   Does knowing the number of each atom in the chemical formula give their mass proportion in the formula? For example, every sample of sulfur trioxide ( ), is approximately 60% sulfur and 40% oxygen, by mass.   This means that a 100 gram sample of sulfur trioxide contains 60 g of sulfur and 40 g of oxygen. Sodium chloride ( NaCl ) is 39.3% Na and 60.7% Cl, by mass. One hundred grams of NaCl contain 39.3 g of sodium and 60.7 g of chlorine. In terms of real-life situations, why is knowing the percent composition of substances important?

5. Calculating the Percent Composition of Compounds The percent composition of a compound can be determined from its formula and the molar masses of the atoms that make it up. There are a few steps to calculating the percent composition of a compound (Let’s practice using H 2 O): 1 mol of H 2 O contains 2 mol of hydrogen and 1 mol of oxygen.   3. Determine the percent composition of water by comparing the masses of hydrogen and oxygen atoms in water to the mass of water. 2. Calculate the mass of the compounds and the constituent atoms. 1 2 1 18 1 16 18 2 16 Assume you have 1 mole of the compound. Therefore, every water molecule contains about 11% of hydrogen and 89% of Oxygen.

6. Worked example TNT (trinitrotoluene) is a white crystalline substance that explodes at 240 °C. Determine the percent composition of TNT, .   STEP 1: Assume you have 1 mol of TNT Calculate the mass of TNT and the atoms in TNT STEP 2: O C H N TNT 1 O C H N TNT 1 6 7 5 3 16 12 1 14 227 96 84 5 42 227

7. O C H N TNT 1 O C H N TNT 1 6 7 5 3 16 12 1 14 227 96 84 5 42 227 STEP 3: Calculate the percentage of each atom in TNT        

8. Activity Calculate the percent composition of aluminum oxide, . What is the percent composition of diphosphorus pentoxide, ? The molecular formula of the insecticide DDT is . Calculate the molar mass of the compound and the percent by mass of each element. Determine the percentage by mass of silver in silver sulphide ( ). How many grams of iron are in 5 kg of ? How many grams of oxygen are in 200 g of ? Chalcopyrite ( ) is a principal mineral of copper. Calculate the number of kilograms of Cu in 3.71 tons of chalcopyrite.  

9. Empirical Formulas List the main types of chemical formulae: Structural formula Molecular formula Empirical formula ? The empirical formula of a compound shows the simplest ratio of ions or atoms in the compound.   Empirical formulae use the smallest possible set of subscript numbers. Molecular formula Empirical formula             Sometimes the molecular and the empirical formulae of a molecular compound are the same, as they are for , , .   Therefore, the molecular formula of a compound can either be the same as or a whole - number multiple of its empirical formula.

10. Determination of Empirical Formulae from Composition Data We can determine the empirical formula of a compound from its composition data. We can determine the mole ratio of each element from the mass to determine the formula of the compound. Consider a 2.476 g sample of an oxide of copper which is found to contain 2.199 g of copper. Solution : Cu O 2.199 0.277 Cu O 2.199 0.277 0.0346 0.0173 3. Determining the empirical formula : Since we conventionally use whole number subscripts, we must simplify the ratio. Divide each number by the smallest number in ratio.   Determine the mass of oxygen in the sample. Determine the moles of Cu and O in the sample. Determine the empirical formula. 1.   2. 63.5 16    

11. Worked example Determine the empirical formula of a compound of phosphorus and oxygen that contains 43.64% phosphorus by mass. Determine the percentage of oxygen in the compound. Assumed you have 100 g of the compound. In 100 g of the compound, there is 43.64 g of phosphorus and 56.36 g of oxygen. Determine the moles of the atoms in the compound.   P O 43.64 56.36 P O 43.64 56.36 1.4077 3.5225 31 16 4. Determine the empirical formula.    

12. Activity A sample of a compound containing boron (B) and hydrogen (H) contains 6.444 g of B and 1.803 g of H. Determine its empirical formula. A 1.500 g sample of a compound contains 0.467 g sulfur and 1.033 g chlorine. What is the empirical formula of the compound? Determine the empirical formula of a compound that is 63.6% nitrogen and 36.4% oxygen. The percentage of carbon, hydrogen and nitrogen in an unknown compound is found to be 23.30%, 4.85% and 40.78% respectively. Determine the empirical formula of the compound. 2.40 g of element Z combines exactly with 1.60 g of oxygen to form a compound with the formula . Determine the relative atomic mass of Z. A sample of a brown-black compound is composed of 2.477 g manganese and 1.323 g oxygen. What is the empirical formula of the compound?  

13. Molecular formula and Empirical formula For each compound: Calculate the molar mass of both the molecular formula and empirical formula, Then divide the molar mass of the molecular formula by the molar mass of the empirical formula. As a general statement: n represents any natural number. n is a factor by which the subscripts of the atoms in the empirical formula are multiplied.   What do you notice ?

14. Exit ticket

15. QUANTITATIVE CHEMISTRY The mole concept

16. PDN 1. What do you think is a molecular formula? 2. Where have you come across the term mole? 3. What was the meaning of this term? 16

17. By the end of the lesson learners should be able to -Define the terms molecular formula and mole. -Determine a molecular formula from percent composition. -State the relationship between the number of elementary particles and the mole. -Calculate the number of particles from the number of moles and the number of moles from the number of particles. 17 Lesson objectives

18. How is the value of n determined?   Worked example The empirical formula of a molecular compound of boron is . The molar mass of a molecule of the boron compound is found to be . Determine the molecular formula of the compound. Solution:     Activity A compound has been isolated and analyzed. The empirical formula of the compound is . Its molecular mass is . What is the molecular formula of the compound? The empirical formula of colorless gas is . Its molecular mass is . What is the molecular formula of the compound? compound used to whiten teeth has the empirical formula of HO. The molecular mass of this compound is . What is its molecular formula? Xylene, a solvent used in industry, has the empirical formula of . The molecular mass of xylene is . What is the molecular formula of xylene?  

19. Introduction to quantitative chemistry What is matter? Quantity  is defined as an amount, measure or number. What do you understand by the word “ Quantity ”? Matter is anything that has mass and takes up space. At a minimum, matter requires at least one subatomic particle , although most matter consist of atoms . Matter can include molecules , compounds, and elements. Physical quantities used to measure matter: Mass : Volume : Mole : is a measure of the amount of matter in a substance or an object. is a measure of the amount of space that a substance or an object takes up. one mole of a substance contains of its elementary particles.   The mole allows scientists to calculate the number of elementary entities (usually atoms or molecules ) in a certain mass of a given substance. Type of substance Elementary particles Molecular substance molecules Ionic substance formula units Element atoms

20. The Mole Concept The mole (abbreviated mol ) is the measure of quantity of the number of elementary particles in a substance. 1 mol of a substance contains  of its elementary particles.   How many elementary particles are in 0.5 mol of a substance? 0.5 mol of a substance contains  of its elementary particles.   How many moles of a substance do elementary particles make up?   elementary particles make up of a substance.     The number of moles a particular number of elementary particles make up is calculated a follows: represents the number of moles of the substance.   represents the number of elementary particles in the substance.   is called the Avogadro’s number     Avogadro’s number  ( N A ) is the numerical value of particles in a mole.  

21. Worked Example 1 How many molecules are present in of ? How many atoms are in this? To determine the total number of atoms, we have to note that there are 3 atoms in a single molecule of water.   EXIT TICKET How many molecules are present in 4.61 × 10 −2  mol of ? How many moles of a substance do   molecules make?  

22. QUANTITATIVE CHEMISTRY Relationship between the mole concept

23. PDN 1. Is there a relationship between the mass of a substance and number of moles? 2. If yes, what is the relationship and if no, give a reason. 23

24. By the end of the lesson learners should be able to -Define the term molar mass. -Calculate the molar mass of any given substance. - State temperature and pressure at STP. -Calculate number of moles from given mass or volume at STP and vice versa. 24 Lesson objectives

25. Molar mass of a specific substance is the mass in grams of one mole of that substance. Molar mass is represented by the symbol .   The unit for molar mass is .   The molar mass of water is .   Worked example Determine the molar mass of . Solution : In , there is 1 Sulphur atom and 3 oxygen atoms. To calculate the molar mass, we will add the molar masses of the atoms in the formula.   Note : the value of the molar mass of a substance is equal to the value of the formula mass of its elementary particle. Activity Which of the following compounds have the greatest molar mass? The herbicide paraquat  ( ) Urea ( The anaesthetic haloethane  ( ) Caffeine ( ) A typical soap,   

26. The relationship between the number of moles in a particular mass of a substance is as follows:   represents number of moles of the substance ( in mol )   represents mass of the substance ( in grams )   is the molar mass of the substance ( in )   Worked example Calculate the number of moles in 50 g of oxygen gas ( ) Calculate the mass of 3.5 moles of . Solution:   Activity Is 6.07 g of equal, greater or less than 1 mole of ? Calculate the number of moles of chloroform ( ) in 198 g of chloroform. Calculate the molar mass of a compound if 0.372 mole of it has a mass of 152 g. How many molecules of ethane ( ) are present in 0.334 g of ? Calculate the number of C, H, and O atoms in 1.50 g of glucose ( ), a sugar.   2.  

27. Moles, mass and number of particles The molar mass of a substance is used to convert between grams of a substance and moles of a substance. How do we convert between moles and mass of a substance? How do we convert between number of particles in a substance and mass of a substance? If we want to convert particles to mass, we must first convert particles to moles and then we can convert moles to mass.     then   OR   leading to   OR  

28. Chapter 9 28 At standard temperature and pressure ( S T P ), 1 mole of any gas occupies (or ).   The volume occupied by 1 mole of gas (22.4 L) is called the molar volume . Molar Volume Standard temperature and pressure are respectively 0  C (273.15 K) and 1 atm (101.3 kPa ). The symbol for molar volume is .   Molar volume has units of .   The moles in a particular volume of gas at STP can be calculated as follows:   represents the moles of the gas at STP (in moles).   represents the volume of the gas at STP ( ).   represents the molar Volume at STP ( ).  

29. Worked Example A sample of methane, CH 4 , occupies at STP. How many moles of methane are present? Solution   EXIT TICKET How many moles of gas occupy at STP? Determine the volume occupied by 0.85 mol of gas at STP. 1.96 g of an unknown gas occupies 1000 at STP. What is the molar mass of the gas? What is the mass of of ozone gas ( at STP? How many molecules of propane gas ( ) occupy at and ?  

30. Calculating moles of atoms in a compound: Calculate the number of nitrogen and oxygen atoms in 1 mole of .   2. Convert the number of nitrogen and oxygen atoms to moles of the atoms. 3. How many moles of N and O are in 1 mole of   Oxygen atoms:   Nitrogen atoms :   There are 2 moles of N and 5 moles of O in 1 mole of .   Activity Determine the number of moles of constituent atoms in 1 mole of the following compounds.   There are molecules of in 1 mol of . In each molecule, there are 2 nitrogen atoms and 5 oxygen atoms. Therefore, in 1 mole of , there are nitrogen atoms, and oxygen atoms.   Note : the subscript numbers in the formula are equal to the number of moles of each atom in one mole of the compound.

31. QUANTITATIVE CHEMISTRY The mole concept and solutions

32. PDN 1. What is a solution? 2. What is the difference between a concentrated solution and a diluted solution? 32

33. By the end of the lesson learners should be able to -Define the term concentration. -Calculate concentration of a solution given mass of solute and volume of solvent . - Calculate number of moles of the solute in a given solution - 33 Lesson objectives

34. Solution: Solutions a homogenous mixture of solute and solvent Solute: the substance that is dissolved in the solution Solvent: the substance in which another substance is dissolved, forming a solution. The solvent is the major component of the solution and the solute is the minor component. Forms of liquid solutions: Aqueous solution: Non-aqueous solution: the solution in which water is the solvent. E.g. salt in water, sugar in water and copper sulfate in water. E.g., iodine in carbon tetrachloride, sulphur in carbon disulfide, phosphorus in ethyl alcohol. the solution in which a solute is dissolved in a solvent other than water. When a gaseous or solid material dissolves in a liquid, the gas or solid material is called the solute. When two liquids dissolve in each other, the major component is called the  solvent   and the minor component is called the  solute . Types of Solutions gaseous solutions ( air, natural gas, LPG, etc ) liquid solutions ( salt water, tea, etc ) solid solutions (brass, bronze, steel, etc ) Why is a solution is considered a mixture? the ratio of solute to solvent can vary. Why is a solution is considered homogeneous ? it has the same composition throughout.

35. Concentration of Solutions The concentration of a solution expresses the amount of solute present in a solution. Concentration: the amount of solute per unit volume of solution.   represents the moles of the solute (in mol )   represents the volume of the solution (in )   represents the concentration of solution (in )     Dilute solution: Concentrated solution: contains a relatively large amount of solute dissolved in a given amount of solution. contains a relatively small amount of solute dissolved in a given amount of solution. NOTE : Acids and bases are purchased from suppliers as concentrated solutions. They are then diluted with water to form dilute solutions.

36. Worked example Calculate the concentration of the solution formed when 4 moles of glucose are dissolved in 5 of water. Solution:   Activity Determine the concentration of 3.0 moles of nitric acid in 4.0 of solution. Calculate the number of moles of hydrochloric acid present in 0.80 of a solution with a concentration of 0.40 . Calculate the concentration of a solution that contains 50.0 g of NaOH in 850 of solution . What volume of a 1.50 copper (II) sulphate solution contains 35.0 g of ? How many grams of NaCl are in 2.53 of a 0.750 solution of NaCl ?   Combining the equations and    

37. Dilution of Solutions Many chemical reagents are supplied in solutions of much higher concentration than their use requires. The concentrated solutions need to be diluted with solvent to prepare solutions of lower concentration. Concentrated solution :   How can a solution be diluted? Adding MORE SOLVENT dilutes a solution State how adding more solvent affects the volume of solution , moles of solute and concentration of solution . the volume of the solution increases without changing the number of moles of solute. the concentration of the solution decreases. Diluted solution :   and are the volume and concentration of the concentrated solution   and are the volume and concentration of the diluted solution   Since the number of moles of solute remains constant with dilution: This yields the dilution equation :   If any three of the four terms are known, the fourth can be calculated. Note that C and V can be in any units provided that the same units are used on both sides of the equation.

38. Worked example Hydrochloric acid is obtained commercially at a concentration of . How many milliliters of the solution of HCl must be used to prepare 2000 of solution of HCl ? Solution :   Activity Calculate the volume to which of hydrochloric acid must be diluted to produce a solution with a concentration of exactly . How would you prepare of a solution of hydrochloric acid starting from a solution? of NaCl is added to of a solution. Calculate the concentration of ions in the final solution.  

39. Chemical Equations Chemical reaction: a process in which one or more substances are changed into one or more new substances. Chemical equation : a standard way to represent and communicate chemical reactions. Types of chemical species in chemical reactions: Reactants : the starting materials in a chemical reaction. Products: the substances formed as a result of a chemical reaction. A chemical equation: shows the formulae (molecular formulae or formula units) of all the reactants and all the products. also gives the number of each species that are required/produced from a complete reaction.

40. EXIT TICKET Activity Write the formulae of the following common compounds: Carbon monoxide Sulphur dioxide Hydrochloric acid Copper (II) sulphate Sulphuric acid Aluminium oxide Magnesium nitrate Iron (III) sulphide Calcium phosphate Ammonium carbonate Nitrous acid Potassium acetate Calcium chlorate Chromium (VI) oxalate Manganese (IV) nitride 2. Write balanced equations for the following reactions. Iron and bromine reacting to form iron (III) bromide. Sodium oxide reacting with sulphuric acid to form sodium sulphate and water. Iron (II) chloride reacting with chlorine to form iron (III) chloride. Sulfur dioxide reacting with hydrogen sulfide to form sulfur and water. Silver reacting with nitric acid to form silver nitrate, nitrogen dioxide and water. Manganese (IV) oxide reacting with hydrochloric acid to form manganese (II) chloride, chlorine and water. Ammonia reacting with oxygen to form nitrogen monoxide and water. Lead(IV) oxide and carbon monoxide forming lead metal and carbon dioxide. Ethanol burning in air to form carbon dioxide and water.  

41. QUANTITATIVE CHEMISTRY Interpretation of chemical equations

42. PDN 1. What is a chemical equation? 2. Are chemical equation important? If yes, given a reason for your answer and if no also given a reason for your answer 42

43. By the end of the lesson learners should be able to State the law of conservation of mass. Interpret a chemical equation in terms of moles, mass and volume. Use chemical equations to calculate unknown quantities. - 43 Lesson objectives

44. Stoichiometry Stoichiometry is the study of quantitative aspects of chemical equations. The amounts of substances in a balanced equation are known as the stoichiometry of the reaction. For example, consider the industrial synthesis of ammonia from hydrogen and nitrogen: The coefficients in balanced equations indicate the number of each species involved in a single chemical reaction. How many of each species are required/produced if reactions are to take place?   Determine the number of moles of each species that are required/produced in these reactions: Therefore, the stoichiometric coefficients in a chemical equation can be interpreted as the ratio of the number of moles of substances in a reaction. In the ammonia synthesis reaction above, one molecule of reacts with three molecules of to form two molecules of .  

45. Quantitative Proof of the Law of Conservation of Mass   Consider the combustion of propane: 1 mol 5 mol 3 mol 4 mol Calculate the mass of each species in the reaction: 3 4 44 32 44 18 44 160 132 72       Therefore, mass is conserved State the law of conservation of mass in words: 1 5 Matter cannot be created or destroyed during a chemical or physical change.

46. Stoichiometry and the Mole Concept the amount of one reactant needed to completely react with a known amount of another reactant. the amount of product that will be formed from specific amounts of reactants. the amount of reactants that must be used to obtain a specific amount of product. The knowledge of stoichiometry is essential in interpreting a reaction quantitatively . The coefficients in a balanced equation can be interpreted as mole ratios of the respective substances. Consider the reaction between methane gas and chlorine gas:   In terms of moles of each substance:   Interpreting a balanced equation in terms of moles allows the possibility of determining

47. Application of Stoichiometry ONLY MOLES are used in STOICHIOMETRY . Worked example When butane is burnt in excess air, the following reaction takes place: Calculate the number moles of oxygen that are required to completely react 5 moles of butane. Solution: The ratio of to is For 5 mol:   Determine the value of x   32.5 moles of oxygen are required to completely react 5 moles of butane.   2. When magnesium is added to aqueous silver nitrate, the following reaction takes place What mass of silver is formed when 3.6 g of magnesium is added to an excess of aqueous silver nitrate? Solution : We must first calculate the moles of magnesium. 0.3 mol of silver is formed form 0.15 mol of magnesium Determining the mass of silver: Therefore, 32.4 g of silver is formed from 3.6 g of magnesium    

48. EXIT TICKET When heated potassium chlorate decomposes to form potassium chloride and oxygen. The unbalanced equation is: Balance the equation. Calculate how many moles of are needed to produce 0.60 moles of oxygen. What mass of is needed to produce 0.02 moles of ? 2. When heated sodium hydrogen carbonate decomposes to form sodium carbonate, carbon dioxide and water. Write a balanced chemical equation for this decomposition reaction. What mass of sodium hydrogen carbonate must be heated to give 8.80 g of carbon dioxide? Calculate the mass of required for the complete combustion of 0.25 mol of propane gas, (g). How many moles of are produced in the complete combustion of 0.75 mole of hexane, ? At high temperatures, aluminum reacts with oxygen to form aluminum oxide. The balanced equation is: How many moles of are consumed for each 0.45 mole of Al? How many moles of will be obtained when 0.45 mole Al is consumed?  

49. QUANTITATIVE CHEMISTRY The limiting reagent

50. PDN What stops you from buying a car? Write down what you think is meant by the term limiting reagent 50

51. By the end of the lesson learners should be able to - Define the term limiting reagent. -Determine a limiting reagent and an excess reagent in a chemical reaction. -Use the limiting reagent to calculate the amount of excess reagent used or unreacted and quantities of products formed. -Use the limiting reagent to determine the percentage yield - 51 Lesson objectives

52. Limiting Reagents When a chemist carries out a reaction, the reactants are usually not present in exact stoichiometric amounts . Stoichiometric amounts mean that the reagents are in the proportions indicated by the balanced equation. The goal of a reaction is to produce the maximum quantity of a useful substance from the starting materials. Frequently, a large excess of one reactant is supplied to ensure that the more expensive reactant is completely converted to the desired product. Consequently, some reactant will be left over at the end of the reaction. The reactant used up first in a reaction is called the limiting reagent (or reactant) When the limiting reactant is used up, no more product can be formed. Excess reagents are the reactants present in quantities greater than necessary to react with the quantity of the limiting reagent. LIMITING REAGENTS determine the AMOUNT OF PRODUCT formed. PHeT Simulation: Limiting reagents

53. In the reaction of sodium hydroxide, NaOH, with hydrochloric acid, HCl ( aq ) , the balanced equation shows that each mole of NaOH reacts with the exact same number of moles of HCl ( aq ) : NaOH and HCl react in a 1 to 1 mole ratio. If NaOH and HCl ( aq ) are combined in any ratio other than a 1 to 1 mole ratio, one reactant will be in excess. Whenever starting amounts of the reactants are given, first determine which reactant is the limiting reactant before calculating amounts of product.  

54. Worked example Consider the reaction: What mass of iodine is produced when 100.00 g of KI is added to a solution containing 12.00 g of and 50.00 g ? Solution: The mole ratio from the equation is The actual mole ratio of reagents present is Simplify the ratio:   Ratio divided by coefficient:   Potassium iodide (KI) is the limiting reagent Amount of iodine produced : Determining mass of iodine : 76.51 g of iodine is produced from 100 g of KI  

55. Activity Consider the reaction: Three starting mixtures of and C are listed below. Determine the limiting reagent in each mixture. 2.0 moles + 4.0 moles C 2.0 moles + 3.0 moles C 2.0 moles + 2.0 moles C 6.00 moles of Al are combined with 6.00 moles of S and heated to form . How many moles of will be produced? 4.00 moles of and 12.0 moles of are reacted according to the reaction How many moles of will form?  

56. Reaction Yield Chemical reactions in the real world don't always go exactly as planned on paper. Factors that could contribute to less product formation: spills and other experimental errors undesirable side reactions reversibility of reactions Chemists need a measurement that indicates how successful a reaction has been. Types of reaction yields theoretical yield : actual yield : the amount of product that would result if all the limiting reagent reacted. the amount of product actually obtained from a reaction.   The actual yield and theoretical yield should be in the same units (e.g. grams, moles, etc ). Low percent yields can cause a large waste of reactants and unnecessary expenses. Thus, chemists strive to maximize the percent yield in a reaction. This measurement is called the percent yield: Predicted amounts of product are sometimes not attained. Yield is a measure of the extent of a reaction, generally measured by comparing the amount of product against the amount of product that is possible.

57. Suppose in a synthesis reaction it was calculated that 50.0 grams of product could be obtained. But when the reaction is performed, only 45.5 grams was obtained. The percent yield would be:   Worked example 0.800 mole of phosphorus was reacted with an excess of sulphur to form diphosphorus pentasulphide , . At the end of the reaction, 82.7 g of was produced. What is the percent yield of this reaction? Solution : First, calculate the moles of that could be obtained from 0.8 mol of phosphorus.   0.4 mol of could be obtained from 0.8 mol of . Determining mass of :  

58. Activity Potassium chlorate decomposes according to the following reaction: In a certain experiment, 40 g is heated until it completely decomposes. What is the theoretical yield of oxygen gas? The experiment is performed and the oxygen gas is collected and its mass is found to be 14.9g . What is the percent yield for this reaction? 2. Mercury(II) oxide decomposes with heating to form elemental mercury, Hg, and oxygen gas, . The decomposition of 0.058 mole of HgO produced 0.724 g of . What is the percent yield of this reaction? 3. When solutions of calcium chloride and silver nitrate are mixed, a precipitate of silver chloride forms. A solution containing 0.32 mole of was treated with an excess amount of . After removing and drying the precipitate, 44.1 g of AgCl was obtained. What is the percent yield of this reaction?