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SCH3 U- R.H. KING ACADEMY MOLE & STOICHIOMETRY HANDOUT Name:

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The Mole- 6.02 x 10 23

ODE TO A MOLE

I find that my heart beat goes out of control Just thinking how useful to man is the mole! So perfectly compact. What could be neater? Only occupying twenty-two and four-tenths of a litre. What kind of equations could one hope to equate Without calculating a formula weight? And could a solution be kept for posterity Without ever knowing its molarity? Though all of these findings may set you to slumber I myself am aroused by Avogadro's number. Life without the mole? Don't be absurd! Count your blessings up to six point zero two times ten to the twenty-third. Mollionaire Q: How long would it take to spend a mole of $1 coins if they were being spent at a rate of 1 billion per

second? A: The Mole A counting unit. Similar to a dozen (12), gross (144), pair (2), and ream (500), except instead of 12, 144, 5 or 500, it is 602 billion trillion, 602,000,000,000,000,000,000,000; 6.02 X 1023 (in scientific notation) This number is named in honor of Amedeo _____________________ (1776 – 1856), who studied quantities of gases and discovered that no matter what the gas was, there were the _____________________number of molecules present.

A Mole of Particles contains 6.02 x 1023 particles

1 mole C = 6.02 x 1023 C atoms 1 mole H2O = 6.02 x 1023 H2O molecules 1 mole NaCl = 6.02 x 1023 NaCl formula units (technically, ionics are compounds not molecules so they are called formula units)

= 6.02 x 1023 Na+ ions and = 6.02 x 1023 Cl– ions

Atomic Mass It is useful to associate atomic mass with a mass in grams. It has been found that: 1 g H-1, 12 g C-12, or 23 g Na-23 have _____________________atoms 6.02 x 1023 is a “mole” or “_____________________” “mol” is used in equations, “mole” is used in writing; one gram = 1 g, one mole = 1 mol. Since the mole is so large, we use it to count very tiny things – like atoms. Because the mole is so large, (and we now know that we cannot count out a mole of anything), how do we know when we have a mole of anything?


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We _____________________ the mass and _____________________ that to the number of atoms present.

Molar Mass The molar mass is the mass of _____ mole of a pure substance. The pure substance can be an element or a compound. The atomic mass is the mass of _____ atom of that element measured in amu’s. The molar mass is also equal to 1 mole of atoms measured in grams.

Comparing sugar (C12H22O11) & H2O

Do they have the same 1 g of each 1 mole of each volume?

mass?

# of moles?

# of molecules?

# of atoms

Atomic Mass vs. Molar Mass

Mass of 1 atom of Pb = ______________amu Mass of 1 mole of Pb atoms = ______________g Mass of 1 atom of N = ______________amu Mass of 1 mole of N atoms = ______________g

Molecular Mass/Molecular Weight (Molecular Compounds): If you have a single molecule, mass is

measured in _____________instead of grams. But, the molecular mass/weight is the _____________

numerical value as 1 mole of molecules. Only the units are different.

Formula Mass/Formula Weight (Ionic Compounds): Same goes for ionic compounds. But again, the

numerical value is the same. Only the units are different.

Molar Mass for Compounds

How to find the molar mass: 1. Write a CORRECT formula for the compound 2. Look up the atomic mass of each element in the compound 3. Multiply the atomic mass by the subscripts, if any. 4. Add all masses of elements together and use the unit, g/mol

Example: Find the molar mass of copper (II) bromate.


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Practice: Calculate molar masses (to 2 decimal places) of:

(NH4)2CO3

O2

Conversions:

The factor- label method (TEXTBOOK, P. 652) To use this we need: 1. desired quantity, 2. given quantity, 3. conversion factor

The steps to follow Now we are ready to solve problems using the factor label method. The steps involved are:

1. Write down the desired quantity/units 2. Equate the desired quantity to given quantity 3. Determine what conversion factors you can use (both universal and question specific) 4. Multiply given quantity by the appropriate conversion factors to eliminate units you don’t want and leave

units you do want 5. Complete the math

Factor-label Method example: Q - How many moles are present in 35.4 grams of Cu?

SCH3 U- R.H. KING ACADEMY MOLE & STOICHIOMETRY HANDOUT Name

H4

Percentage Composition Q: If you have a box containing 100 golf balls and 100 Ping-Pong balls, which type of ball contributes the most to the mass of the box? A: The same principle applies to finding the % composition of a compound. Different elements have different masses and this must be taken into consideration.

Percent Composition (by mass)

Percent Composition (by mass): Identifies the elements present in a compound as a mass percent of the total compound mass. The mass percent is obtained by ____________________the mass of each element by the______________ Mass of a compound and ____________________ to percentage.

% composition of an element =mass of elementmass of compound x 100%

How to find the percent composition of a compound: 1. Write a correct formula for the compound 2. Find the molar mass of the compound 3. Divide the total atomic mass of EACH ELEMENT by the molar mass 4. Multiply by 100 to convert your results to a percent 5. Since you have no significant figures to go by, express your answer to TWO decimal places with the % sign. Example: What is the percent composition of each element in NH4OH by mass? Practice: 1. Find the percentage composition by mass of aluminum thiocyanate. (Al- 13.41%, S- 47.80%, C-17.91%, N- 20.88%) 2. A student prepares a compound of tungsten chloride from 3.946 g of tungsten and 3.806 g of chlorine. Assuming the reaction goes to completion, calculate the percent composition. (W- 50.50%, Cl-49.50%) 3. How many grams of sodium will combine with 567.0 g of sulfur to form Na2S? (Total Mass= 1379.56 g, Mass of Na= 812.56g)


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Empirical Formula The ____________________ ratio of elements in a compound. It uses the ____________________ possible ____________________number ratio of atoms present in a formula unit of a compound. If the percent composition is known, an empirical formula can be calculated.

Simplest and molecular formulae

Chemical formulas are either “simplest” (a.k.a. “empirical”) or “molecular”. Ionic compounds are ________________

expressed as ____________________formulas. Covalent compounds can either be molecular formulas (i.e. ______________) or simplest (e.g.

__________________)

Q - Write simplest formulas for propene (C3H6), C2H2, glucose (C6H12O6), octane (C8H14) A: Q - Identify these as simplest formula, molecular formula, or both H2O, C4H10, CH, NaCl. A: How to find the empirical formulae of a compound: Calculations to find the simplest formula incorporate this rhyme: % to mass Mass to mole Divide by small Multiply till whole A chart form may help to organize work. Example 1: A compound contained 29.08 % Na, 40.58 % S, and 30.34 % O. Find the empirical formula of this compound.

species mass or % molar mass 𝒏 =𝐦𝐚𝐬𝐬 𝐨𝐫 %𝐦𝐨𝐥𝐚𝐫 𝐦𝐚𝐬𝐬

divide by smallest from

column 4

reduce to smallest whole number ratio

Compound Formula Empirical Formula

Hydrogen peroxide Benzene Ethylene Propane


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Example 2: A 5.72 g sample of washing soda (Na2CO3 . x H2O) is heated to give 2.12 g of anhydrous Na2CO3.

What is the simplest formula of the hydrated salt? species mass or % molar mass 𝒏 =

𝐦𝐚𝐬𝐬 𝐨𝐫 %𝐦𝐨𝐥𝐚𝐫 𝐦𝐚𝐬𝐬


column 4


Example 3: A compound contained 40.0g C, 6.71g H, and 53.3g O. Find its empirical formula and the empirical formula mass.



column 4


Empirical Formula Mass =

Mole ratios and simplest formula: Given the following mole ratios for the hypothetical compound AxBy, what would x and y be if the mole ratio of A and B were:

Formula: A = 1 mol, B = 2.98 mol A = 1.337 mol, B = 1 mol A = 2.34 mol, B = 1 mol A = 1 mol, B = 1.48 mol If any result from Step 3 is a mixed number, you must multiply ALL values by some number to make it a whole number. Ex: 1.33 x 3; 2.25 x 4; 2.50 x 2, etc. Formulas for Compounds Empirical Formula ________________ possible set of subscript numbers Smallest ________________ number ratio All ionic compounds are given as _____________ formulas

Molecular Formula The ________________ formulas of molecules It shows __________ of the atoms present in a molecule It may be the ________________ as the E.F. or a whole number________________ of its E.F.


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n represents a ________________number multiplier from 1 to as large as necessary How to find the molecular formulae of a compound: 1. Calculate the empirical formula and the mass of the empirical formula 2. Divide the given molecular mass by the calculated empirical mass 3. Answer is a whole number multiplier 4. Multiply the empirical formula by the multiplier Example: Lactic acid has a molar mass of 90.08 g and has this percent composition: 40.0% C, 6.71% H, 53.3% O. What is the empirical and molecular formula of lactic acid? Assume a 100.0 g sample size. Step 1. Use Chart to find the Empirical Formula:



column 4


Step 2. Obtain the mass of the Empirical Formula:

Step 3. Obtain the value of n (whole number multiplier):

Step 4. Multiply the empirical formula by the multiplier Questions: 1. What information must be known to determine a) the empirical formula of a substance? b) the molecular formula of a substance? 2. Determine the molecular formula for each compound below from the information listed. substance simplest formula molar mass(g/mol) naphthalene C5H4 128

)/()/(molgmassformulaempirical

molgmassmolarn=


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Part 1: Stoichiometry (Mass to Moles)

2 Mg(s) + O2(g) → 2 MgO(s)

Note that the Law of Conservation of Mass is always obeyed. Consider: 4NH3 + 5O2 → 6H2O + 4NO In words, this tells us that for every _______ moles of NH3, _______ moles of O2 are required Is 4 g NH3 / 5 g O2 a conversion factor?

Stoichiometry questions (1) -Factor-Label Method Consider: 4NH3 + 5O2 → 6H2O + 4NO

1. How many moles of H2O are produced if 0.176 mol of O2 are used?

2. How many moles of NO are produced in the reaction if 17.00 mol of H2O are also produced?

3. How many grams of H2O are produced if 1.9 mol of NH3 are combined with excess oxygen?

4. How many grams of O2 are required to produce 0.3 mol of H2O?

5. How many grams of NO is produced if 12 g of O2 is combined with excess ammonia?


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Moving along the stoichiometry path We always use the same type of information to make the jumps between steps:

grams (x) ↔ moles (x) ↔ moles (y) ↔ grams (y)

Example:

Consider: 4NH3 + 5O2 → 6H2O + 4NO a) How many moles of H2O can be made using 0.50 mol NH3?

b) What mass of NH3 is needed to make 1.50 mol NO?

c) How many grams of NO can be made from 120.0 g of NH3? The steps to solving stoichiometric problems are as follows- CHART METHOD 1. Write the chemical equation 2. Balance the chemical equation 3. Write the molar ratio for the equation 4. Write all given masses of substances in the equation 5. Write the molar masses for all substances 6. Find the number of moles of each substance 7. Find the new molar ratio 8. Solve for the unknown Example: Given 4NH3 + 5O2 → 6H2O + 4NO How many grams of NO can be made from 120.0 g of NH3? 4 NH3 5 O2 6 H2O 4 NO Molar Ratio(MR)

Mass (m)

Molar Mass (MM)

Moles (n)


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Stoichiometry- Limiting Reagents A balanced chemical equation indicates the number of ________________of each REACTANT that will react and the number of ________________of each PRODUCT that will be produced in the reaction Even if one reactant is present in excess only the amount ________________to react, dictated by the molar ratio, will actually react. The amount reacting will be determined by the reactant that is in the ________________amount (_________________). Problems of this type may be recognized by the fact that information will be given about at ________________two reactants. Q - How many moles of NO are produced if __ mol NH3 are burned in __ mol O2? 4 mol NH3, 5 mol O2 4 mol NH3, 20 mol O2 8 mol NH3, 20 mol O2

In limiting reagent questions we use the limiting reagent as the “given quantity” and ________________ the reagent that is in excess. Example: How many grams of NO are produced if 4 moles NH3 are burned in 20 mol O2?

Solving Limiting Reagents: grams to moles Q - How many g NO are produced if 20 g NH3 is burned in 30 g O2?

1. First we need to calculate the number of moles of each reactant NH3 O2 Molar Ratio

What we Have

Limiting Factor Test

The _______________value indicates the Limiting Reactant Determine the reactant in excess and the excess amount if 20 g NH3 is burned in 30 g O2? Reactant NH3 O2

Molar Ratio

What we have

Limiting Factor Test

What we need

Excess


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Stoichiometry (given = limiting) 1) Express all chemical quantities as moles 2) Determine the limiting reagent via a chart 3) Perform the stoichiometry using the limiting reagent as the “given” quantity Example: How many g NO are produced if 20 g NH3 is burned in 30 g O2? Given: 4NH3 + 5O2 → 6H2O + 4NO 4NH3 5O2 6H2O 4NO Molar Ratio (MR) Mass (m)

Molar Mass (MM)

Moles (n)

PERCENTAGE YIELD (P. 336-338)

Yield:

Theoretical yield:

Actual yield: Q: Give 4 possible reasons why the actual yield in a chemical reaction often falls short of the theoretical yield.

Percent Yield =


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Sample problem:

What is the % yield of H2O if 138 g H2O is produced from 16 g H2 and excess O2? [143 g; 96.7%]

Step 1: Write the balanced chemical equation:

Step 2: Determine actual and theoretical yield. Actual is given, theoretical is calculated:

Step 3: Calculate % yield:

Practice problem: What is the % yield of NH3 if 40.5 g NH3 is produced from 20.0 mol H2 and excess N2? [227 g, 17.8%]

Practice problem: When 5.00 g of KClO3 is heated it decomposes according to the equation: 2KClO3 → 2KCl + 3O2

a) Calculate the theoretical yield of oxygen. [1.96 g]

b) Give the % yield if 1.78 g of O2 is produced. [90.9%]

c) How much O2 would be produced if the percentage yield was 78.5%? [1.537 g O2]

the mole- 6.02 x 10 23 - wikispacessch3uking.wikispaces.com/file/view/mole... · we _____ the mass...

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