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**Chemical Quantities or**

Chapter 10 Chemical Quantities or "Our Friend the Mole"

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**How you measure how much?**

You can measure mass, or volume, or you can count pieces. We measure mass in grams. We measure volume in liters. We count pieces in MOLES.

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**Moles 1 mole is 6.02 x 1023 particles.**

Particles can be atoms, molecules, formula units. 6.02 x is called Avogadro's number.

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**Representative particles**

The smallest pieces of a substance. For an element it is an atom. Unless it is diatomic For a molecular compound it is a molecule. For an ionic compound it is a formula unit.

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**Conversion factors Used to change units. Three questions**

What unit do you want to get rid of? Where does it go to cancel out? What can you change it into?

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Calculation question How many molecules of CO2 are the in 4.56 moles of CO2 ? 4.56 moles CO2 x 6.02 x 1023 Molecules = 2.75 x 1024 Molecules CO2 1 Mole CO2

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**Calculation question How many moles of water is 5.87 x 1022 molecules?**

5.87 x 1022 molecules H2O x 1 Mole H2O____________ = Moles H2O 6.02 x 1023 Molecules H2O

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Calculation question How many atoms of carbon are there in 1.23 moles of C6H12O6 ? 1.23 moles C6H12O6 x 6 moles C Atoms_____ x 6.02 x 1023 Atoms C 1 mole C6H12O6 Molecules mole C atoms = 4.44 x atoms C

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**Measuring Moles The decimal number on the periodic table is**

The mass of the average atom in amu The mass of 1 mole of those atoms in grams.

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Gram Atomic Mass Gram Atomic Mass is the mass of 1 mole of an element in grams. We can write this as g C = 1 mole We can count things by weighing them.

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**Examples How much would 2.34 moles of carbon weigh?**

2.34 moles C x 12 g C__ = 28.1 g C 1 mole C

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**Examples How many moles of magnesium in 4.61 g of Mg?**

4.61 g Mg x 1 mole Mg = moles Mg 24.3 g Mg

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**Examples How many atoms of lithium in 1.00 g of Li?**

1.00 g Li x 1 mole Li x x 1023 atoms Li = 8.7 x atoms Li 6.9 g Li mole Li

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**Examples How much would 3.45 x 1022 atoms of U weigh?**

3.45 x 1022 atoms U x 1 mole U_________ x 238g U = g U 6.02 x 1023 atoms U mole U

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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. (Be careful) To find the mass of one mole of a compound Determine the moles of the elements contained in the compound. Find out how much each would weigh add them up

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What about compounds? Molar Mass is the mass of one mole of any compound. Example: What is the Molar mass of CH4? 1 mole of C = g 4 mole of H = 4 x 1.01 g = 4.04g 1 mole CH4 = = 16.05g/mol

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**Molar Mass What is the molar mass of Fe2O3?**

2 moles of Fe = 2 x g = g 3 moles of O = 3 x g = g The Molar mass = g g = g/mol

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**Calculate the molar mass of the following**

C6H12O6 C = 6 x 12g = 72g H = 12 x 1.0g = 12g O = 6 x 16g = 96g Molar mass = = 180g/mol (NH4)3PO4 N = 3 x 14g = 42g P = 1 x 30.9g = 30.9g O = 4 x 16g = 64g Molar mass = 42g + 12g g + 64g = 148.9g/mol

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**Finding moles of compounds Counting pieces by weighing**

Using Molar Mass Finding moles of compounds Counting pieces by weighing

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Molar Mass The number of grams in 1 mole of atoms, formula units, or molecules. We can make conversion factors from these. To change grams of a compound to moles of a compound. Or moles to grams

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

5.69 g NaOH x 1 mole NaOH Molar mass of NaOH Molar Mass NaOH = Na + O + H = 23.0g+16g + 1.0g = 40g/mole Substitute 40g/mole into the equation above and get: 5.69 g NaOH x 1 mole NaOH = 0.14 moles NaOH 40 g NaOH

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Gases and the Mole

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

They are difficult to weigh, so we’ll measure volume Need to know how many moles of gas we have. Two things affect the volume of a gas Temperature and pressure Compare at the same temp. and pressure.

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**Standard Temperature and Pressure**

Avogadro's Hypothesis - at the same temperature and pressure equal volumes of gas have the same number of particles. 0ºC and 1 atmosphere pressure Abbreviated atm 273 K and kPa kPa is kiloPascal

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**At Standard Temperature and Pressure**

abbreviated STP At STP 1 mole of gas occupies 22.4 L Called the molar volume Used for conversion factors Moles to Liter (x 22.4), and L to mol (divide by 22.4)

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

4.59 mole CO2 x 22.4L CO = L CO2 1 mole CO2

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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 than the mass is the molar mass (PT) At STP the volume is 22.4 L.

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Examples Find the density of CO2 at STP. Note: At STP the volume is 22.4 L D =m/v = Molar mass CO2 = ((12g + (16g x 2g)) = 1.96g/l 22.4 L L

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**Quizdom Find the density of CH4 at STP.**

D=m/v ((12.0g +(1.0g x 4)) = g/l 22.4L

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The other way Given the density, we can find the molar mass of the gas. Again, pretend you have a 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. What is the molar mass of a gas with a density of g/L? M=D x V = 1.964g/l x 22.4L = 44g

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**All the things we can change**

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**Mass Volume Moles 6.02 x 1023 Representative Particles Ions Atoms**

PT Moles Representative Particles 6.02 x 1023 Ions Atoms Count

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**Percent Composition Like all percents Part x 100 % whole**

Find the mass of each component, divide by the total mass.

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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. Part x 100 Whole Total weight = 29.0g Ag g S = 33.3 g %Ag = 29.0g Ag x 100 = 87.1% Ag 33.3 g % S = 4.30g S x 100 = 12.9% S 33.3g

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**Getting it from the formula**

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

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**Examples Calculate the percent composition of C2H4? Assume 1 mole**

1 mole C2H4 = (12g x 2) +(1.0g x 4) = 28g/mole total mass % C = (12g x 2) x = 85.7% C 28 g % H = (1.0g x 4) x = 14.3% H

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**Examples What is the percent composition of Aluminum carbonate.**

Write the formula: Al2(CO)3 (ionic Formula writing) Assume 1 mole Calculate Total Mass: (Al x 2) + (C x 3) + (O x 3) = (27g x 2) + (12g x 3) + (16g x 3) = 54g + 36g + 48g = 138g % Al = (27g x 2) x = 39.1%Al 138g % C = (12g x 3) x = 26.1% C % O = (16g x 3) x = % O

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Percent to Mass Multiply % by the total mass to find the mass of that component. How much aluminum in 450 g of aluminum carbonate? Write the formula: Al2(CO)3 (ionic Formula writing) % Al = (27g x 2) x = 39.1%Al 138g % C = (12g x 3) x = 26.1% C % O = (16g x 3) x = % O 39.1%Al x1/100 x 450g Al2(CO)3 = 176g Al

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**From percentage to formula**

Empirical Formula From percentage to formula

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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 empirical formula C2H4 molecular formula C3H6 molecular formula H2O both

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**Finding Empirical Formulas**

Just find the lowest whole number ratio C6H12O6 CH4N2 It is not just the ratio of atoms, it is also the ratio of moles of atoms.

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**Calculating Empirical Formulas**

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

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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 = 16.1 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 3.220 molN 1 mol N**

The ratio is mol H = 5 mol H molN mol N C1H5N1 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**

Caffeine is 49.48% C, 5.15% H, 28.87% N and 16.49% O. What is its empirical formula? Since the empirical formula is the lowest ratio the actual molecule would weigh the same or more. By a whole number multiple. Divide the actual molar mass by the the mass of one mole of the empirical formula. You will get a whole number. Multiply the empirical formula by this.

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Example A compound has an empirical formula of ClCH2 and a molar mass of g/mol. What is its molecular formula? A compound has an empirical formula of CH2O and a molar mass of g/mol. What is its molecular formula?

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**Percent to molecular Take the percent x the molar mass**

This gives you mass in one mole of the compound Change this to moles You will get whole numbers These are the subscripts Caffeine is 49.48% C, 5.15% H, 28.87% N and 16.49% O. It has a molar mass of 194 g. What is its molecular formula?

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Example Ibuprofen is % C, 8.80 % H, % O, and has a molar mass of about g/mol. What is its molecular formula?

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1 Chapter 10 “Chemical Quantities” Yes, you will need a calculator for this chapter!

1 Chapter 10 “Chemical Quantities” Yes, you will need a calculator for this chapter!

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