Relate 1 mol of C3H8 is stoichiometrically equivalent to
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Roessler Notes Chemistry Stoichiometry Notes Stoichiometry - branch of chemistry that deals with the relative quantities of reactants and products in chemical reactions - It is a quantitive description of the proportions by moles of the substances in a chemical reaction - The stoichiometry of a reaction is the description of the relative quantities by moles Step Back - The mole concept is the biggest stepping stone for students to overcome - unlike other units that you are used to,l you can not physically see the mole. - It is just like any other unit that we use by counting or weighing - It is used to measure small particles - Its along the same line as a dozen of eggs, a box of pencils, etc - Mole is used by chemist to count chemical entities by weighing them - We will weigh them against a set standard- just like any other unit - Concept of the Mole - The Mole (mol) is the SI unit for the amount of substance - The amount of substance that contains the same number of entities as the number of atoms in 12 grams of Carbon-12 - Also known as Avogadro’s number- to honor his work in chemistry - 1 mol= 6.022E23 entities (I would program this into your calculator - is a unit of measurement for the amount of substance or chemical amount - it is one of the base units in the International System of Units )learned earlier) - it is used to describe chemical reactions History of the Mole - The name mole is a translation in 1897 from the Germans- Mol - Coined by Wilhelm Ostwald in 1893 106751343 1 Roessler Notes Chemistry - It is derived from the German word for molecules- Molekul - A mole is defined as the amount of substance that contains as many elementary entities in it - atoms - molecules - ions - electrons - For a long time we have been using Carbon-12 as our reference tool - Recently that has come into question - Under this concept 1 mole of pure Carbon 12 has a mass of exactly 12 g - The number of entities (atoms, molecules, ions, electrons) in one mole is called Avogadro’s number - Avogadro’s number is a proportionality facto that relates molar mass of an entity to the mass of another entity - It is a constant of 6.02 E 23 - Named after Italian Amedeo Avogadro - Proposed it in 1811 - He proposed that the volume of a gas (at given pressure and temperature) is proportional to the number of atoms or molecules regardless of how the gas acts - French physicist Jean Perrin in 1909 proposed naming the constant after Avogadro - Jean Perrin would go on later to win a Nobel Prize for his work on deriving the Avogadro’s constant through different methods - The value of the constant was proposed by Johann Josef Loschmidt - 1865 he estimated the average diameter of the molecules in airthe method is similar to calculating the number of particles in a given volume of gas 106751343 2 Roessler Notes Chemistry - Accurate determinations would come from two famous scientist - This would require the measurement of a single quantity on the atomic and macroscopic scale using the same unit of measurement - Robert Millikan did this by measuring the charge on an electron in 1910 - Michael Faraday discovered the charge of one mole of electrons (he is well known for his work on electrolysis) - electrolysis- chemical decomposition produced by passing an electric current through a liquid or solution containing ions - by dividing the charge on a mole of electrons by the charge of a single electron the value that we know today as 6.02 E 23 was derived - Since 1910 we have developed newer methods and newer equipment that better allows us to find the exact number - However the number that was discovered then was extremely close to the one today - Perrin proposed the name Avogardo’s Number (N) to refer to the number of molecules in one gram molecule of oxygen - Avogadro’s number is symbolized with Na- it is a base unit in the International System of Units - It is recognized as the amount of substance and is an independent dimension of measurement - It is a unit of measurement - Defining the Mole - 1 mole of carbon-12 contains 6.022 E23 atoms and has a mass of 12 grams - Check out the periodic table! What is the atomic mass unit of Carbon? I wonder if there is a connection between the 12 grams and the AMU of 12 - so if you measured out only 6 grams of Carbon-12, then that represents 0.5 mols of carbon-12 and contains only 3.011 E 23 atoms 106751343 3 Roessler Notes Chemistry - Knowing the amount (in moles), the mass (in grams), and the number of entities becomes very important when we begin to mix substances together - The central relationship between masses on the atomic scale and on the macroscopic scale is the same for elements and compounds - Elements - The mass in atomic mass units (amu) of one atom of an element is the same numerically as the mass in grams (g) of 1 mole of atoms of the element - 1 atom of Sulfur- has a mass of 32.07 amu and 1 mol (6.02E23 atoms) of S has a mass of 32.07 grams - 1 atom of Fe has a mass of 55.85 amu and 1 mol (6.02E23 atoms) of Fe has a mass of 55.85 grams - Note also that since atomic masses are relative - 1 Fe atom weighs 55.85/32.07 as much as 1 S atom - 1 mol of Fe weighs 55.85/32.07 as much as 1 mol of S - Compounds - The mass in atomic mass units (amu) of one molecule (or formula unit) of a compound is the same numerically as the mass in grams (g) of 1 mole of the compound - 1 Molecule of water- has a mass of 18.02 amu and 1 mol (6.02E23 atoms) of water has a mass of 18.02grams - 1 formula unit of NaCl has a mass of 58.44 amu and 1 mol (6.02E23 atoms) of Fe has a mass of 58.44 grams - Note again that because masses are relative, 1 water molecule weighs 18.02/58.44 as much as 1 Sodium Chloride formula unit, and 1 mol of water weighs 18.02/58.44 as much as 1 mol of NaCl - The two key points to remember about the importance of the mole unit - The mole lets us relate the number of entities to the mass of a sample of those entities - The mole maintains the same numerical relationship between mass on the atomic scale (atomic mass unit, AMU) and mass on the macroscopic scale (grams, g) - Determining Molar Mass 106751343 4 Roessler Notes Chemistry - The Molar mas (M) of a substance is the mass per mole of its entities (atoms, molecules, or formula units) and has units of grams per mole (g/mol). You will be using the periodic table to do this - 1. elements - to find the molar mass, look up the atomic mass and note whether it is a monoatomic or molecular - Remember that there are seven diatomic and they make a 7 with the exception of one who is always the exception - Monatomic elements- the molar mass is the periodic table value in grams per mole - Example Gold is 197.0 g/mol - Example Neon is 20.18 g/mol - The mass value in the periodic table has no units because it is a relative atomic mass, given by the atomic mass (in amu) divided by 1 amu (1/12 mass of one Carbon-12 atom in amu) - Relative atomic mass= atomic mass (amu)/1/12 mass of Carbon 12 (amu) - Therefore, you use the same number for the atomic mass and for the molar mass - Molecular elements- you must know the formula to determine the molar mass. - For example in air, oxygen exists most commonly as diatomic molecules, so the molar mass of Oxygen is twice that of Oxygen - Molar mass of O2= 2 times 16.00= 32.00 g/mol - Sulfur can exist as an octatomic molecule - Its mass is 8 times 32.07 g/mol = 256.6 g/mol - 2. Compounds- The molar mass is the sum of the molar masses of the atoms in the formula - Sulfur Dioixde= S (32.07 g/mol) + 2 O (16.00 *2)= 64.07 g/mol 106751343 5 Roessler Notes Chemistry - Potassium Sulfide (K2S)--> 2*K (2* 39.10 g/mol) + S (32.07)= 110.27 g/mol - This the subscripts in a formula refer to individual atoms (or ions) as well as to moles of atoms (or ions) - Uses of Stoichiometry - It prevents accidents in the lab and it prevents waste - we are able to predict how much we need and how much it will produce - It can be used to calculate quantities such as the amount of products (in mass, moles, volume, etc) that can be produced with given reactants and percent yield - percent yield- the percentage of the given reactant that is made into the product - It can predict how elements and components diluted in a standard solution react in experimental conditions - It is founded on the law of conservation of mass - the mass of the reactants equal the mass of the products - Types of Stoichiometry - Reaction stoichiometry - describes the quantitive relationships among elements in compounds - EX describes the 1:3:2 ratio of molecules of nitrogen, hydrogen, and ammonia - Composition stoichiometry - describes the quantitive (mass) relationships among elements in compounds - Ex composition stoichiometry describes the nitrogen to hydrogen relationship for the production of ammonia: 1 mole of nitrogen and three moles of hydrogen are in every mole in ammonia - gas stoichiometry - deals with reactions involving gases; where the gases are at a known temperature, pressure, and volume - We assume that all the gases are behaving as ideal gases 106751343 6 Roessler Notes Chemistry - this is a theoretical gas that is composed of randomly moving non interacting particles - since they do not interact we are able to examine the way gases behave under specific temperature and pressure - Though we generally set the temperature and pressure to be at STP - At low temperatures or higher pressures the ideal gas model tends to fail- at these points the intermolecular forces and molecular size becomes important - It also fails when dealing with heavy gases such as water vapor and many refrigerants - Newtonian and Quantum Physics have been used to explore these gases relationships - Also Maxwell-Boltzmann concept is considered to be the classical concept of ideal gases - Balancing Equations - Based on the law of conservation of mass (and energy) - Neither created nor destroyed but transferred - Mass is independent of gravity or the attraction back to earth - inertial mass - Weight is not independent of gravity rather it is mass times gravity - Basically it will remain the same or constant over time as long as the system is closed - closed system- it is an isolated system that cannot exchange heat, work, or matter with the surroundings - Open system- can exchange all of heat, work, and matter - The law was discovered by Antoine Lavosier in the late 18th century - changed alchemy into chemistry which is why he is the father of chemistry - Basically “Nothing comes from Nothing” 106751343 7 Roessler Notes Chemistry - statement from Empedocles - Joesph Black, henry Cavendish, and Jean Rey also helped to develop ideas to support the law •Calculating Quantities of Reactant and Products •A balanced equation is essential for all calculations involving chemical changes •If you know the number of moles of one substance, the balanced equation tells you the number of moles of the other •Stoichiometrically Equivalent Molar Ratios • the amounts (mol) of substance are stoichiometrically equivalent to each other •Which means that a specific amount of the other •the quantitative relationships are expressed as stoichiometrically equivalent molar ratios that we use as conversion factors to calculate the amounts •Lets look at a reaction Propane (C3H8) and oxygen react and undergo complete combustion •If we view the reaction quantitatively in terms of propane we see that In relationship to C3H8 • Therefore in this reaction 106751343 8 Roessler Notes Chemistry Relate 1 mol of C3H8 is stoichiometrically equivalent to • We could have also looked at any other element in this reaction in the same way • A balanced equations contains a wealth of quantitative information relating individual chemical entities, amounts (mols) of substances, and masses of substances Amount (mol) 1 mol C3H8 5 mol O2 --> 3 mol CO2 4 mol H2O Molecules 1 molecule C3H8 5 molecules of O2 --> 3 molecules of CO2 4 molecules of H2O 106751343 9 Roessler Notes Chemistry Mass (amu) 44.09 amu C3H8 160.00 amu O2 --> 132.03 amu CO2 72.06 amu H2O mass (g) 44.09 g C3H8 160.00 g O2 --> 132.03 g CO2 72.06 g H2O Total mass (g) 204.09 g --> 204.09 g • The coefficient represents the number of moles in the reaction for each reactant and product • Balanced equations always represent moles and not grams you get grams by calculating the formula mass or molecular masses • If I had 10.0 mol of H2O and I wanted to find out the amount of Oxygen that was needed to produce it then I would do the following 106751343 10 Roessler Notes Chemistry 10.0 mol H2O converted to Oxygen • You cannot solve this type of problem without the balanced equation • Here is an approach for solving any stoichiometry problem that involves a reaction • 1. Write the balanced equation • 2. When necessary, convert the known mass (or number of entities) of one substance to amount (mol) using its molar mass (or Avogadro’s number) • 3. Use the molar ratio to calculate the unknown amount (mol) of the other substance • 4. When necessary, convert the amount of that other substance to the desired mass (or number of entities) using its molar mass (or Avogadro’s number) Homework Problems a. In a lifetime, the average American uses 1750 lb (794 kg) of copper in coins, plumbing, and wiring. Copper is obtained from sulfide ores, such as chalocite (copper (I) sulfide) by a multistep process. After grinding the ore, it is roasted (heated strongly with oxygen gas) to form powdered copper (I) oxide and gaseous Sulfur dioxide. How many moles of Oxygen are required to roast 10.0 mol of copper (I) Sulfide 106751343 11 Roessler Notes Chemistry • b. During the roasting process, how many grams of sulfur dioxide form when 10.o mol of copper (I) sulfide reacts • c. During the roasting of chalcocite, how many kilograms of oxygen are required to form 2.86 kg of Copper (I) oxide • Reactions that occur in a Sequence • In many situations, a product of one reaction becomes a reactant for the next in a sequence of reactions • when the same (common) substance forms in one reaction and reacts in the next, we eliminate it in an overall (net) equation • Those that occur in one reaction and react in the next are called spectator ionsthey are not our major focus- so we eliminate • The steps to writing the overall reaction are • 1. Write the sequence of balanced equations • 2. Adjust the equations arithmetically to cancel the common substance • 3. Add the adjusted equations together to obtain the overall balanced equations Example Problem 106751343 12 Roessler Notes Chemistry Roasting is the first step in extracting copper from chalcocite. In the next step copper (I) oxide reacts with powdered Carbon to yield Copper metal and carbon monoxide gas. Write a balanced overall equation for the two step process. • Reactions that involve a Limiting Reactant • So far we have looked at the amount of one reactant that was given. • We assumed there was enough of the other reactants to react with it completely Copper (I) sulfide reacts with Oxygen to form Copper(i) oxide and Sulfur Dioxide • We assume that 5.2 mols of Cu2S reacts with as much Oxygen as needed • Because all the Cu2S reacts the initial amount of 5.2 mol determines or limits the amount of SO2 that can form, no matter how much more Oxygen is present • We therefore call the Copper (I) sulfide the limiting reagent • Suppose however you know the amounts of both Copper (I) sulfide and Oxygen and need to find out how mush Sulfur dioxide forms • The reactant that is not limiting is present in excess • Excess means that you have leftover- you had too much in the reaction • To determine which is the limiting reactant, we use molar ratios in the balanced equation to perform a series of calculations to see which reactant forms less products 106751343 13 Roessler Notes Chemistry • Determining the limiting reagent • you will be given the quantities of two or more reactants • then you will have to determine the limiting reagent based on molar ratios • you must have a balanced equation • you are looking for the one that produces the least amount of products Example Problem Nuclear engineers use chlorine trifluoride to prepare uranium fuel for power plants. The compound is formed as a gas by the reaction of elemental chlorine and fluorine. (Recall that Halogens have distinct colors as you move down the groups and generally get darker... Remember also the Chlorine was used in WWII as mustard gas) HWK 106751343 14 Roessler Notes Chemistry 1. In another preparation of Chlorine trifluoride, 0.750 mol of Cl2 reacts with 3.00 mol of F2. Find the limiting reagent 2. A fuel Mixture used in the early days of rocketry consisted of two liquids, hydrazine (N2H4) and dinitrogen tetraoxide which ignites on contact to form nitrogen gas and water vapor. How many grams of nitrogen gas forms when 1.00 E 2 g of N2H4 and 2.00 E 2 g of N2O4 are mixed • Theoretical, Actual, and Percent Reaction Yield • we have assumed that 100% of the limiting reagent becomes product 106751343 15 Roessler Notes Chemistry • That is if you have a perfect lab • In reality it doesn’t always work that way • Theoretical Yield • The amount of product calculated from the molar ratio in the balanced equations • There are several reasons why this is never obtained • Reactant mixtures often proceed through side reactions that form different products- these will decrease the amount available to react • Many reactions stop before they are complete • Physical losses occur in every step of separation; some solid clings to filter paper, some distillate evaporates, and so forth- think of your lab mistakes • Actual yield • The amount of product that is actually yielded • Theoretical and actual amounts are expressed in units of amounts (moles) or mass (grams) • Percent yield • Is the actual yield expressed as a percentage of the the theoretical yield • ratio of actual to theoretical and then multiplied by 100 for percent Percent Yield • Practice Problem 106751343 16 Roessler Notes Chemistry Silicon carbide is an important ceramic material made by reacting sand (silicon dioxide) with powdered carbon at high temperature. Carbon monoxide is also formed. When 100.0 kg sand is processed, 51.4 kg od Silicon carbide is recovered. What is the percent yield of Silcon carbide in this process. • Fundamentals of Solution Stoichiometry • liquid solutions are easier to store than gases and easier to mix than solids, and the amounts of substances in solution can be measured precisely • Many environmental and biological reactions occur in solutions • Expressing Concentration in terms of Molarity • a solution consists of a smaller quantity of one substance- the solute • It is dissolved in the solvent • When it dissolves, the solute’s chemical entities become evenly dispersed throughout the solvent • Water is considered to be the universal solvent • It dissolves a lot of things because of its bonds! It is covalent, but acts like an ionic because of the electrons on the molecule. • The concentration of a solution is often expressed as the quantity of solute dissolved in a given quantity of solution • Concentration is an intensive property (like density or temperature) and thus independent of the solution volume 106751343 17 Roessler Notes Chemistry • Molarity (M)- expresses the concentration in units of moles of solute per liter of solution Molarity (M) Example Problem Glycine has the simplest structure of the 20 amino acids that make up proteins. What is the Molarity of the solution that contains 0.715 mol of glycine in 495 mL? Amount- Mass- Number Conversions Involving Solutions • Like many intensive properties, Molarity can be used as a conversion factor between Volume (L) of solution and amount (mol) of solute, from which we can find the mass or the number of entities of solute 106751343 18 Roessler Notes Chemistry Amount- Mass- Number Example Problem Biochemist often study reactions in solutions containing phosphate ion, commonly found in cells. How many grams of solute are in 1.75 L of 0.460 M sodium hydrogen phosphate? • Preparing and Diluting Molar Solutions • The volume term in the denominator of the molarity expression is the solution volume- not the solvent volume • This is because the solute volume adds to the solvent volume, the total volume (solute + solvent) would be more than 1 L, so the concentration would be less than 1M • Preparing s Solution • there are four steps to do this properly- Lets prepare 0.500 L of 0.350 M Nickel(II) nitrate hexahydrate 106751343 19 Roessler Notes Chemistry • 1. weigh the solid- calculate the mass of the solid needed by converting from volume (L) to amount (mol) and then to mass (g) • Transfer the solid- choose the appropriate size container and then add enough distilled water- remember I showed you that tap water can have adverse effects due to the ions in the solution. Make sure you wash down the solid that gets caught on the neck of the volumetric flask • Dissolve the solid- swirl the flask until all the solute is dissolved- you can also preheat the water to help with the dissolving • Add Solvent to the final volume- get up to the line that is etched on the flask. Cap and invert until it gets fully dissolved • Diluting a Solution • a concentrated solution (higher molarity) is converted to a dilute solution (lower molarity) by adding solvent • this means the solution volume increases but the amount (mol) of solute stays the same • The dilute solution contains fewer solute particles per unit volume and thus has a lower concentration than the concentrated solution • you can prepare these from stock solutions- this is what I do for you. Example Problem 106751343 20 Roessler Notes Chemistry Isotonic saline is 0.15 M aqueous Sodium Chloride. It simulates the total concentration of ions in many cellular fluids, and its uses range from cleaning contact lenses to washing red blood cells. How would you prepare 0.80 L of isotonic saline from 6.0 M stock solution • Solving Dilution Problems • To solve dilution problems and others involving change in concentration Formula For Dilutions • M and V are the molarity and volume of the dilute and concentration • Use the values in the problem from above and solve for volume of the concentration Using the equation 106751343 21 Roessler Notes Chemistry • Stoichiometry of Reactions in Solution • Balance the equation • Find the amount (mol) of one substance from the volume and molarity • Relate it to the stoichiometrically equivalent amount of another substance • Convert to the desired units 106751343 22
