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The MOLE Concept Lets talk about it!

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NOT THIS KIND OF MOLE!

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**The mole is used mostly in chemistry as a SI unit for amount of substance.**

In simplest terms a mole represents a number of objects. Suppose you’re engaged in a building project: laying a floor. You’re going to put down 25 boards with 10 nails a piece: that’s at least 250 nails you’ll need. When you go to the hardware store and ask for 250 nails, the clerk doesn’t count out the nails for you. What do they do?

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You weigh them!! When we need a large number of a small item, it doesn’t make sense to count them out; weighing is more efficient. We probably won’t get exactly the number we want, but if we’re dealing in large numbers it won’t make much difference if there are a few more or less.

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In chemistry, we’re stuck with using just this method to count atoms and molecules, because it’s the only way we can do it. Q: Can we count out atoms one-by-one? A: No, they are too small, but we do need to know how many we have in a bunch so we can study their behavior, which is what chemistry is all about. Q: In chemistry labs or when cooking or baking at home, how do we measure out chemicals/reactants/ingredients? A: We either measure it out per a certain volume amount, such as in a graduated cylinder or in a measuring cup, or we weigh it out, like on a balance.

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Why is this important? It is important to know how much of a particular substance we have in order to study its behaviour. Weighing atoms in order to count atoms is not an exact science. If we were to weigh out 250 nails it would be less accurate than actually counting each nail. However, it is much more practical to weigh the nails.

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Because atoms and molecules are so small, chemists created a quantity called a mole that represents a large number of atoms. The unit of a mole is like a dozen in that it represents a certain number of atoms or molecules. What is a dozen? What do we call 2 things?

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**When can we use the terms a dozen or a pair**

When can we use the terms a dozen or a pair? We use terms such as dozen or pair in order to represent a particular number of objects.

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So what is a mole??? It is a very, very, very large number but the same idea as a pair, a dozen, a gross, and so on. A pair is 2 items, a dozen is 12 items, a gross is 144 items, and a mole is 6.02 × 1023 items. That is 602,000,000,000,000,000,000,000 items in a mole. This number is usually expressed in scientific notation.

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**One mole of anything is equal to 6.02 x1023 of those things...**

1 mol of elephants = elephants 1 mole of water molecules =

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A long while a go a man by the last name of Avogadro came up with an idea of how to group atoms! The mole is also referred to as Avogadro’s number! If two elements were at equal volume, temperatures and pressure, he figured there would be the same number of particles in each.

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**But how big is a mole? How big is a mole of basketballs?**

What is another name for the mole? Why do we use the mole?

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Q: can a dozen represent any number besides 12 or can a pair represent a number other than two? A: No, these numbers are constants, which mean that they always represent the exact same number regardless of the item. The mole is most specifically used to count atoms. Even though a mole is like a dozen, the actual number is much larger because atoms are much smaller than the eggs or oranges we usually measure by the dozen.

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**Converting Moles to Particles and Particles to Moles**

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**Recall Dimensional Analysis**

Suppose you buy three and a half dozen roses and you want to know how many roses you have. You can multiply the known quantity (3.5 dozen) by a conversion factor to express the quantity in the units you want (number of roses).

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**Conversion factor: 12 roses 1 dozen **

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**Now lets deal with particles:**

NOTE: 1 MOLE = 6.02 X 1023 PARTICLES PARTICLES (ATOMS OR MOLECULES) How many particles of sucrose are in 3.50 moles of sucrose?

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**Calculate the number of moles that contain 4.50 x 1024 atoms of zinc.**

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**Example 3: How many molecules are in 4. 00 mols of NaCl. 4**

Example 3: How many molecules are in 4.00 mols of NaCl? 4.00 mols NaCl x 6.02 x 1023molecules 1 mol NaCl = 2.41 x 1024 molecules NaCl

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**Questions: How is a mole similar to a dozen?**

Why do chemists use the mole? Do the following conversions using dimensional analysis: How many atoms in 2.50 mol Zn? How many moles contain 5.75 x 1024 atom Al? How many moles contain 3.58 x 1023 formula units ZnCl2 ? How many formula units given 3.25 mol AgNO3 ? How many molecules in 11.5 mol H2O? How many moles contain 3.75 x 1024 molecules CO2 ? How many moles contain 2.50 x 1020 atoms Fe?.

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Question of the Day: Explain how you can convert from the number of representative particles of a substance to moles of that substance.

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Mass of a Mole Avogadro’s number is also used to represent how many atomic mass units can be found in a gram. This allows us to use the mass numbers on the periodic table for both the mass of an atom (atomic mass) and the mass of a mole of atoms (molar mass), we only need to change the units. What is atomic mass? Atomic mass is the weighted average mass of the isotopes of that element. Molar mass – the mass in grams of one mole of any pure substance.

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**Calculating Molar Mass**

Atomic mass = molar mass amu = g/mol What is the molar mass of copper? What is the molar mass of carbon?

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**How can the molar mass of copper and carbon be different if a mole is always 6.02x 102 3 atoms?**

This is the same as asking how the weight of a dozen elephants can be different than the weight of a dozen ants. It is the number of items that is always the same, not the mass of weight or size of those items.

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**What is the molar mass of H2O? What is the molar mass of Ca(OH)2 ? **

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**Calculate the molar mass of the following:**

a) N2O5 b) H2CO3 c) NaC2H3O2 d) CuClO3 e) Ba(NO2)3 f) Pb(HSO4)2

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**Grams Moles, Moles Grams**

Imagine that you are buying jellybeans. Instead of counting out each jellybean they are measured by mass. You find that 1 dozen has a mass of 35 g. What mass of jellybeans should you measure if you want 5 dozen? 35g jellybeans / 1 dozen

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Example 1: Suppose that while working in a chemistry lab you need 3.00 moles of manganese for a chemical reaction. How can you measure that amount?

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Example 2: How many moles of calcium are in 525 g calcium?

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**How many moles of NaCl are in 234.0g NaCl?**

How many grams would 2 moles of NaCl be?

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**Practice Determine the mass in grams of each:**

3.57 mol Al b) 42.6 mol Si c) 3.45 mol H2O d) 2.45 mol ZnCl2 Determine the number of moles in each: 25.5 g Ag b) 125g Zn c) 300.0g Na2S d) 1.00 kg Fe

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**Mass to Atoms, Atoms to Mass Conversions**

Just as we could not make a direct conversion from cm km we cannot make a direct conversion from mass to atoms. Instead we use two step dimensional analysis. Mass Moles Atoms Atoms Moles Mass

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Examples: How many atoms are in 25.0 g of gold?

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**A party balloon contains 5. 50 x 1022 atoms of He gas**

A party balloon contains 5.50 x 1022 atoms of He gas. What is the mass in grams of the helium?

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**How many atoms are in each of the following samples?**

0.230 g Pb 11.5g Hg 45.6 g Si 0.120kg CaCl2 What is the mass in grams of each of the following: a) 1.00 x 1024 atoms Mn b) 3.40 x 1022 atoms He c) 1.50 x 1015 atoms LiCl d) 1.50 x 1015 atoms U

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Hand – In Questions: How many atoms are in 55.2 g Li? What is the mass of 6.02 x 1024 atom Bi? How many molecules are in g of calcium iodide? What is the mass of 3.34 x 1022 molecules H2O? How many sodium chloride atoms in 117g of NaCl?

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**Volume of Gases Determining the Number of particles in a Gas**

Although the size can differ greatly between different gases Avogadro's principle states that all gases that have the same number of particles will have the same volume atoms of Xenon will have the same volume as Helium. The molar volume for gases is the volume that 1 mole of gas will take up at 0 Celsius at 1.00 atm. Since all gases have the same volume at STP (standard temperature and pressure), this means that 1 mol of any gas will have a volume of 22.4L. ** Conversion Factor: 1 mole = 22.4L at STP STP = Standard temperature and pressure (temp = 0oC, pressure = 1 atm)

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**Moles to Volume and Volume to Moles**

Example1: What is the volume of a balloon that holds 1.78 moles of oxygen gas at STP? Example 2: How many moles of helium at STP will fit into a 39.6 L container?

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**Volume to Particles and Particles to Volume**

Example 1: How many particles are in 3.72 L of gas at STP? Example 2: What volume is occupied by 3.24 x 1024 molecules of H2O gas at STP?

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**Volume to Mass and Mass to Volume**

Example 1: Determine the mass of 42.1L of NO2(g) @ STP. Example 2: Determine the volume of the 1.00g of nitrogen gas at STP.

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**Welcome to Stoichiometry**

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Background Comes from the Greek words stoicheion, meaning element and metron, meaning measure. In the real world, chemists perform reactions to make useful chemical compounds such as medicines, fibers, pigments, cosmetics and structural materials. Because some reactants are expensive chemists don’t just mix huge quantities of chemicals together and hope for the best; rather they use stoichiometry to calculate the mass relationships between products and reactants.

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Terminology: Stoichiometry: The Study of relationships between the amounts of reactants used and products formed by a chemical reaction, based on the Law of Conservation of Mass. mole ratio: The ratio of the product and reactant in a chemical reaction. The numbers that fit into this ratio are pulled directly out of the balanced equation for the reaction.

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**Mole ratios are the key to stoichiometric calculations.**

They give us the ratios between the amounts of reactants and products in any given reaction. We get mole ratios from the coefficients. That’s why we need to start with a balanced equation.

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Chemical Equation: 2 Al(s) + 3Br2(l) 2AlBr3(s) For every 2 moles of aluminum, 3 moles of bromine are used to produce the reaction. ** The Law of Conservation is important when dealing with stoichiometry. If the equation isn’t properly balanced then you will use the incorrect mole ratios in your calculations.

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Mole Ratios Example: Write all possible mole ratios 2 Al(s) + 3Br2(l) 2AlBr3(s)

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Quizam Next Wednesday What is a mole? How big is a mole? Who is Avogadro? Molar/ Molecular Mass Dimensional Analysis - particles, volume, mass, moles Law of Conservation of Mass Steps to solve a stoichiometric equation. (Write a balanced chemical equation, convert to moles, use the mole ratio, and find what required.)

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**Determine All Possible Mole Ratios**

Make sure to BALANCE! Al (s) + O2 (g) Al2O3(s) Fe (s) + H2O (l) Fe3O4 (s) + H2 (g) HgO (s) Hg(l) + O2 (g)

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**Mole to Mole Conversions**

Moles of known x moles of unknown = moles of unknown moles of known

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Example 1: How many moles of H2 are we going to need to react with 12.0 moles of O2 to produce H2O2? How much H2O2 will this produce? Step 1: Write balanced equation H2 + O2 H2O2 Step 2: Write down what you are given and convert to moles. Step 3: Use mole ratios to convert to desired chemical.

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**Example2: How many moles of hydrogen are produced when 0**

Example2: How many moles of hydrogen are produced when mole of potassium is used? 2K(s) + 2H2O(l) 2KOH(aq) + H2(g)

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**Particles to Particles Conversions**

Particles moles Particles **Always convert to moles first! 2K(s) + 2H2O(l) 2KOH(aq) + H2(g) How many atoms of H2 are produced if you start with 3.02 x 1022 atoms of H2O?

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Grams to Grams a) Determine the mass of sodium chloride produced when 1.25 grams of chlorine gas reacts with sodium. 2 Na (s) + Cl2 (g) 2NaCl (s)

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Volume to Volume b) The volume of chlorine needed in order to produce 26.0 L of NaCl.

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**Can you convert grams to volume?**

c) Determine the volume of chlorine gas used when 1.25 grams of sodium chloride is produced.

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Practice: 2NaN3 (s) 2Na(s) + 3N2 (g) a) Determine all mole ratios. b) Determine the mass of N2 produced if g of NaN3 is decomposed. c) Determine the volume of N2 gas produced if 2.4 x1023 atoms of NaN3 are used. d) How many moles of sodium are produced if 7.2 mol of NaN3 reacts? e) What must you always convert to first?

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Chapter 11. 1. Describe how a mole is used in chemistry. 2. Relate a mole to common counting units. 3. Convert the number of moles to the number.

Chapter 11. 1. Describe how a mole is used in chemistry. 2. Relate a mole to common counting units. 3. Convert the number of moles to the number.

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