The Mole and Stoichiometry Mr. Kinton Enloe High School Honors Chemistry

Units of Measurement and the Metric System Scientific work makes observations on qualitative and quantitative data Some quantities that chemists measure include: Density Mass Volume Energy

Units of Measurement and the Metric System To help communicate more clearly across the world SI units were introduced Preferred metric units for use in science All conversions are done by either multiplying or dividing by some multiple of 10 Prefixes are used to indicate the size of the unit Without units and a common system, it would be more difficult for scientists to communicate their findings

The Metric System

Common Units in Chemistry Grams (g) and kilograms (kg) Meter (m) and nanometer (nm) Joule (J), calorie (cal), and electron volt (eV) Seconds (s) and minutes (min) Liter (L), milliliter (mL), and microliter (μL) Mole (mol) Kelvin (K) and Celsius (oC) Pascals (Pa), atmospheres (atm), millimeters of mercury (mmHg) What quantities are being measured with these units?

Unit Conversion Given the following quantities convert them to the desired unit: 22.5 mg to g 15.0 mL to μL 10.5 g to kg 15 GB to MB

Scientific Notation A way to represent large or small numbers using less digits Example: 4,000 = 4 × 103 0.000005 = 5 × 10-6 6,325,000 = 6.325 × 106

Recording Measurements and Significant Digits Chemistry is a laboratory science and data is collected based on laboratory experiments: How does a Chemist know when they have included enough decimals in their measurement? Significant figures: the digits that indicate the precision with which a measurement is made All digits of a measured quantity are significant but the last digit is uncertain

Significant Figures We will examine several measurements: 1.350 miles 39.10 grams 810. kJ In order to determine how many significant figures are included in your measurement check the interval on your measuring device Lets examine 2 problems in your packet (numbers 1 and 3)

Counting Significant Digits in a Measurement When in lab we have to look at the instrument we are using but in calculations we will need to examine the measured numbers given in our problems Here are the rules that must be followed when determining how many significant figures a measurement contains: All numbers between 1-9 are significant If you have zeros then follow these guidelines

Determining Which Zeros are Significant Zeros between nonzero digits are always significant Zeros at the beginning of a number are never significant; they are indicating the position of the decimal point Zeros that are at the end of a number AND come after the decimal point are always significant When a number ends in zeros but contains no decimal point, the zeros may or may not be significant

Determining the Number of Significant Figures Determine how many significant figures each of the following measurements contains: 87009 km 0.0095897 m 85.00 g 7,000.00 cm 7.000 × 103 cm

Calculations and Significant Figures When doing calculations you can only be as precise as your least precise measurement Instead of just choosing a location to round off to at random we use the significant figures of our measurements guide our rounding Additionally, when recording measurements, your final answer should only contain one uncertain digit

Multiplication and Division with Sig Figs Answers with multiplication and division must be reported with the same number of significant figures as the measurement with the fewest significant figures 6.221 cm × 5.2 cm = 32.3492 cm2 rounds to 32 cm2 23.0 cm × 432 cm × 19 cm = 188,784 cm3 rounds to 190,000 cm3 In rounding your number, look at the leftmost digit to be dropped: If the leftmost digit to be removed is less than 5 the preceding number is unchanged If the leftmost number to be removed is 5 or greater, increase the preceding number by 1

Addition and Subtraction with Sig Figs The result can have no more decimal places than the measurement with the fewest number of decimal places 20.4 cm + 1.322 cm + 83 cm = 104.722 cm or 105 cm 123.25 mL + 46.0 mL + 86.267 mL = 255.507 mL or 255.5 mL Rounding rules are the same as with multiplication and division

Dimensional Analysis/Factor Label Method Method of problem solving in which the units are carried through all calculations Ensures that the final answer has the correct units Gives a systematic way to solve problems and to help identify possible errors Conversion factor: a fraction whose numerator and denominator are the same quantity expressed in different units

Dimensional Analysis/FLM If an object is 8.50 in long, how many cm long is the object? If a woman has a mass of 115 lbs, what is her mass in grams?

The Mole and Particles The amount of matter that contains as many particles as the number of atoms in exactly 12g of Carbon-12 Avogadro’s number: the number of atoms present in a 12g sample of Carbon-12 which is 6.02×1023 1 mol Carbon-12 atoms = 6.02×1023 carbon-12 atoms 1 mol H2O molecules = 6.02×1023 water molecules 1 mol NO3- ions = 6.02×1023 nitrate ions Using the mole allows us to move from the atomic scale to the macroscopic scale

The Mole and Particles How many moles of magnesium are in 3.01×1022 atoms of magnesium? How many molecules are there in 4.00 moles of glucose, C6H12O6?

The Mole and Molar Mass A mole will always correspond to 6.02×1023 , however the mass of a mole of 2 different objects will be different This allows us to make the connection that 1 atom of Carbon-12 has a mass of 12 amu At the laboratory level, 1 mol of Carbon-12 has a mass of 12g

The Mole and Molar Mass This relationship also works when we examine the formula weights of compounds as well 1 molecule H2O has a mass of 18.02 amu 1 mol H2O has a mass of 18.02g Mass in grams of 1 mol of a substance is the molar mass

The Mole and Molar Mass Give the molar mass for the following compounds: KCl Ca(NO3)2 CuSO4 ● 5H2O How many moles are in 28g of CO2? What is the mass of 5.0 mol of Fe2O3?

The Mole and Volume We can also use the mole to help us determine the volume of a gas at standard temperature and pressure (0o C and 1 atm) 1 mol of any gas = 22.4 L With the mole we can convert between particles, mass, and volume.

The Mole and Volume Determine the volume, in liters, occupied by 2.0 moles of a gas at STP. How many moles of argon are present in 11.2 L of argon gas at STP?

Conversions between Moles While it is useful to convert to moles sometimes we will want to move to another unit Follow this process: Given unit  Moles  Desired Unit How many glucose molecules are in 5.23g of C6H12O6? How many O atoms are in a sample of 4.20g of HNO3?

Percent Composition When given a formula we know the number of atoms or ions present in that compound Percent composition by mass will allow us to determine the empirical formula of a compound Empirical formula: compound which shows the relative number of atoms in a compound

Percent Composition There are 2 ways to determine the percent composition of a compound: % Composition = (subscript of element × molar mass of element) ÷ molar mass of compound % Composition = mass of component in a sample ÷ total mass of the sample Anytime you determine the formula of a compound using percent composition it will be an empirical formula

Percent Composition Ascorbic acid contains 40.92% C, 4.58% H, and 54.50% O by mass. What is the empirical formula of ascorbic acid? A 5.325g sample of methyl benzoate, a compound used in the manufacture of perfumes is found to contain 3.758g of carbon, 0.316g of hydrogen, and 1.251g of oxygen. What is the empirical formula?

Molecular Formula Tells us the actual number of atoms of each element in the compound Will always be a whole number multiple of the empirical formula In order to determine the molecular formula, you must know the molecular weight of your compound This is done by comparing the empirical formula weight to the molecular formula weight

Molecular Formula Step 1: Find the empirical formula Step 2: Calculate the formula weight of the empirical formula Step 3: Find the molecular weight given in the problem Step 4: Divide the molecular weight by the formula weight. Round to a whole number Step 5: Multiply all subscripts in the empirical formula by the whole number

Molecular Formula Earlier we determined the empirical formula of ascorbic acid to be C3H4O3. The experimentally determined mass of this compound is 176 amu. Find the molecular formula. Ethylene glycol, the substance used in antifreeze, is composed of 38.7% C, 9.7% H, and 51.6% O by mass. Its molar mass is 62.1 g/mol. What is the molecular formula for ethylene glycol?

Stoichiometry Quantitative study of reactants and products in a chemical reaction As a result, you must have a balanced chemical equation in order to perform a stoichiometry calculation The coefficients in the equation tell you how many moles of each substance must combine to form products In the reaction: 2H2 (g) + O2 (g)  2H2O (l) 2 moles of hydrogen reacts with 1 mole of oxygen to form 2 mole of liquid water If we can balance an equation, we can determine a myriad of chemical quantities based on that equation.

Stoichiometry As we solve problems using stoichiometry follow these steps: Write the correct formulas for all reactants and products. Balance the equation Convert the quantities of the given or known substances into moles Use the mole ratio (coefficients) from the balanced equation to convert to the moles of the substance needed Convert the moles needed to the unit that is specified in the problem Check that you answer is reasonable

Stoichiometry How many grams of water are produced in the oxidation of 1.00 g of glucose? Calculate the volume of CO2 produced by the combustion of 1.00 g of butane.

Stoichiometry and Limiting Reactants During most reactions a reaction will stop once a reactant is completely used up Limiting reactant/reagent: the reactant that is completely consumed during the course of a chemical reaction Excess reactant/reagent: reactant that is present in a greater quantity and is left over once the reaction stops

Limiting Reactant Urea is prepared by reacting ammonia and carbon dioxide. Water is the other product. Consider the reaction of 637.2 g of NH3 with 1,142 g of CO2. What is the maximum mass of urea that can be formed? What is the limiting reactant? How much of the excess reactant remains?

Limiting Reactant Consider the reaction between aluminum and iron (III) oxide. In one process, 124 g of aluminum reacts with 601 g of iron (III) oxide. What is the maximum mass of iron that can be produced? What is the limiting reactant? How much of the excess reactant will remain once the reaction has finished?

Reaction Yield Anytime a reaction takes place it will yield a certain amount that is less than what we would expect This is due to side reactions that take place and our inability to capture all products Reaction yield or percentage yield is the amount of product expressed as a percentage of the actual yield and the theoretical yield of the products

Reaction Yield Actual Yield: the amount of product that is experimentally produced Theoretical Yield: the calculated amount from the balanced equation. This will be the amount of product formed based on the limiting reactant % Yield = (Actual yield ÷ Theoretical yield) × 100

Reaction Yield In a certain industrial operation, 3.54 × 107 g of TiCl4 reacted with 1.13 × 107 g of magnesium. Calculate the theoretical yield of Ti Calculate the percent yield of Ti if 7.91 × 106 g were obtained in the operation

Reaction Yield Consider the reaction of vanadium (V) oxide. Calculate the theoretical yield of vanadium in a process that involves the reaction of 1.54 × 103 g of vanadium (V) oxide with 1.96 × 103 g of calcium. What is the percent yield if 803 g of vanadium were obtained in the process?