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**Section 7.1 The Mole: A Measurement of Matter**

OBJECTIVES: Describe how Avogadro’s number is related to a mole of any substance. Calculate the mass of a mole of any substance.

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**What is a Mole? You can measure mass, or volume,**

or you can count pieces. We measure mass in grams. We measure volume in liters. We can count pieces by counting atoms, formula units or molecules.

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**Moles (abbreviated: mol)**

Defined as the number of carbon atoms in exactly 12 grams of carbon-12 (C-12 an isotope that has mass of 12). 1 mole is 6.02 x particles. Treat it like a very large dozen 6.02 x is called Avogadro’s number.

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**Representative particles**

The smallest pieces of a substance. For a molecular compound: it is the molecule. For an ionic compound: it is the formula unit (ions). For an element: it is the atom. Remember the 7 diatomic elements (made of molecules)

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**Types of questions How many oxygen atoms in the following? CaCO3 3**

Al2(SO4)3 12 How many ions in the following? CaCl2 NaOH 2 5

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**Types of questions (Particles)**

How many moles of water is 5.87 x molecules? How many moles is 7.78 x 1024 formula units of MgCl2? Units MUST be the same so they can cancel out. Units for final answer. 5.87 x 1022 molecules 1 mole = mol 1 6.02 x 1023 molecules 7.78 x 1024 formula units 1 mole = 12.9 mol 1 6.02 x 1023 formula units

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**Types of Questions (to Moles)**

How many molecules of CO2 are there in 4.56 moles of CO2? 4.56 mol 6.02 x 1023 molecules = 2.75 x 1024 molecules 1 1 mole

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**Measuring Moles Remember relative atomic mass?**

The amu was one twelfth the mass of a carbon-12 atom. Since the mole is the number of atoms in 12 grams of carbon-12, the decimal number on the periodic table is also the mass of 1 mole of those atoms in grams.

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Gram Atomic Mass (gam) Equals the mass of 1 mole of an element in grams 12.01 grams of C has the same number of pieces as grams of H and grams of iron. We can write this as g C = 1 mole C We can count things by massing them.

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**Examples How much would 1.00 moles of carbon weigh? 12.01 grams**

How many moles of magnesium is g of Mg? 1 mol How many atoms of lithium is 1.00 g of Li? 8.67 x 1022 atoms How much would 3.45 x 1022 atoms of U weigh? 13.64 grams

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What about compounds? in 1 mole of H2O molecules there are two moles of H atoms and 1 mole of O atoms To find the mass of one mole of a compound determine the moles of the elements they have Find out how much they would weigh add them up

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**What about compounds? What is the mass of one mole of CH4?**

1 mole of C = g 4 mole of H x 1.01 g = 4.04g 1 mole CH4 = = 16.05g The Gram Molecular Mass (gmm) of CH4 is 16.05g this is the mass of one mole of a molecular compound.

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**Gram Formula Mass (gfm)**

The mass of one mole of an ionic compound. Calculated the same way as gmm. What is the GFM of Fe2O3? 2 moles of Fe x g = g 3 moles of O x g = g The GFM = g g = g

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Note! Gram Atomic Mass, Gram Molecular Mass, Gram Formula Mass are all generally called Molar Mass. As it sounds, molar mass is the mass of exactly 1 mole (in grams) of that substance.

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**Section 7.2 Mole-Mass and Mole-Volume Relationships**

OBJECTIVES: Use the molar mass to convert between mass and moles of a substance.

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**Section 7.2 Mole-Mass and Mole-Volume Relationships**

OBJECTIVES: Use the mole to convert among measurements of mass, volume, and number of particles.

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Molar Mass Molar mass is the generic term for the mass of one mole of any substance (in grams) The same as: 1) gram molecular mass, 2) gram formula mass, and 3) gram atomic mass- just a much broader term.

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Examples Calculate the molar mass of the following and tell what type it is: Na2S N2O4 C Ca(NO3)2 C6H12O6 (NH4)3PO4 78.04 g/mol ionic compound 92.02 g/mol molecular compound 12.01 g/mol atom 164.1 g/mol ionic compound g/mol molecular compound g/mol ionic compound

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**Molar Mass The number of grams of 1 mole of atoms, ions, or molecules.**

We can make conversion factors from these. To change grams of a compound to moles of a compound.

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For example How many moles is 5.69 g of NaOH?

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For example How many moles is 5.69 g of NaOH?

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**For example How many moles is 5.69 g of NaOH?**

need to change grams to moles

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**For example How many moles is 5.69 g of NaOH?**

need to change grams to moles for NaOH

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**For example How many moles is 5.69 g of NaOH?**

need to change grams to moles for NaOH 1mole Na = 22.99g 1 mol O = g 1 mole of H = 1.01 g

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**For example How many moles is 5.69 g of NaOH?**

need to change grams to moles for NaOH 1mole Na = 22.99g 1 mol O = g 1 mole of H = 1.01 g 1 mole NaOH = g

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**For example How many moles is 5.69 g of NaOH?**

need to change grams to moles for NaOH 1mole Na = 22.99g 1 mol O = g 1 mole of H = 1.01 g 1 mole NaOH = g

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**For example How many moles is 5.69 g of NaOH?**

need to change grams to moles for NaOH 1mole Na = 22.99g 1 mol O = g 1 mole of H = 1.01 g 1 mole NaOH = g

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**Examples How many moles is 4.56 g of CO2? 0.1 mol**

How many grams is 9.87 moles of H2O? g How many molecules is 6.8 g of CH4? 2.55 x 1023 molecules 49 molecules of C6H12O6 weighs how much? 1.47 x grams

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**Gases Many of the chemicals we deal with are gases.**

They are difficult to weigh. Need to know how many moles of gas we have. Two things effect the volume of a gas Temperature and pressure We need to compare them at the same temperature and pressure.

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**Standard Temperature and Pressure**

0ºC and 1 atm pressure abbreviated STP At STP 1 mole of gas occupies 22.4 L Called the molar volume 1 mole = 22.4 L of any gas at STP

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**Types of Questions (Volume)**

How many moles are 3.97 L of STP 1 mole of ALL gases take up 22.4 L of space)? Remember: STP Standard Temperature and Pressure (0°C or 273K temperature and 101.3kPa or 1 atmosphere pressure). 3.97 L 1 mole = mol 1 22.4 L

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**Examples What is the volume of 4.59 mole of CO2 gas at STP?**

L of CO2 How many moles is L of O2 at STP? 0.25 mol of O2 What is the volume of 8.8 g of CH4 gas at STP? L of CH4

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**Density of a gas D = m / V for a gas the units will be g / L**

We can determine the density of any gas at STP if we know its formula. To find the density we need the mass and the volume. If you assume you have 1 mole, then the mass is the molar mass (from PT) At STP the volume is 22.4 L.

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**Examples Find the density of CO2 at STP. 1.965 g/L of CO2**

Find the density of CH4 at STP. 0.716 g/L of CH4

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The other way Given the density, we can find the molar mass of the gas. Again, pretend you have 1 mole at STP, so V = 22.4 L. m = D x V m is the mass of 1 mole, since you have 22.4 L of the stuff.

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The other way What is the molar mass of a gas with a density of g/L? g/mol 2.86 g/L? g/mol

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**Summary These four items are all equal: a) 1 mole**

b) molar mass (in grams) c) 6.02 x 1023 representative particles d) 22.4 L at STP Thus, we can make conversion factors from them.

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**Section 7.3 Percent Composition and Chemical Formulas**

OBJECTIVES: Calculate the percent composition of a substance from its chemical formula or experimental data.

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**Section 7.3 Percent Composition and Chemical Formulas**

OBJECTIVES: Derive the empirical formula and the molecular formula of a compound from experimental data.

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**Calculating Percent Composition of a Compound**

Like all percent problems: Part whole Find the mass of each component, then divide by the total mass. x 100 %

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Example Calculate the percent composition of a compound that is 29.0 g of Ag with 4.30 g of S.

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**Getting it from the formula**

If we know the formula, assume you have 1 mole. Then you know the mass of the pieces and the whole.

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**Examples Calculate the percent composittion of C2H4?**

How about Aluminum carbonate? Sample Problem 7-11, p.191 We can also use the percent as a conversion factor Sample Problem 7-12, p.191

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The Empirical Formula The lowest whole number ratio of elements in a compound. The molecular formula = the actual ratio of elements in a compound. The two can be the same. CH2 is an empirical formula C2H4 is a molecular formula C3H6 is a molecular formula H2O is both empirical & molecular

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**Calculating Empirical**

Just find the lowest whole number ratio C6H12O6 CH4N It is not just the ratio of atoms, it is also the ratio of moles of atoms. In 1 mole of CO2 there is 1 mole of carbon and 2 moles of oxygen. In one molecule of CO2 there is 1 atom of C and 2 atoms of O.

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**Calculating Empirical**

We can get a ratio from the percent composition. Assume you have a 100 g. The percentages become grams. Convert grams to moles. Find lowest whole number ratio by dividing by the smallest.

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Example Calculate the empirical formula of a compound composed of % C, % H, and %N. Assume 100 g so 38.67 g C x 1mol C = mole C gC 16.22 g H x 1mol H = mole H gH 45.11 g N x 1mol N = mole N gN

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**Example The ratio is 3.220 mol C = 1 mol C 3.219 molN 1 mol N**

The ratio is mol H = 5 mol H molN mol N = C1H5N1 A compound is % P and % O. What is the empirical formula? Caffeine is 49.48% C, 5.15% H, 28.87% N and 16.49% O. What is its empirical formula?

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**Empirical to molecular**

Since the empirical formula is the lowest ratio, the actual molecule would weigh more. By a whole number multiple. Divide the actual molar mass by the empirical formula mass. Caffeine has a molar mass of 194 g. what is its molecular formula?

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Example A compound is known to be composed of % Cl, 24.27% C and 4.07% H. Its molar mass is known (from gas density) to be g. What is its molecular formula? C2H4Cl2 Sample Problem 7-14, p.194

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