Download presentation

Presentation is loading. Please wait.

2
**Chemical Quantities or**

Chapter 10 Chemical Quantities or "Our Friend the Mole"

3
**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.

4
**Moles 1 mole is 6.02 x 1023 particles.**

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

5
**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.

6
**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?

7
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

8
**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

9
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

10
**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.

11
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.

12
**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

13
**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

14
**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

15
**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

16
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

17
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

18
**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

19
**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

20
**Finding moles of compounds Counting pieces by weighing**

Using Molar Mass Finding moles of compounds Counting pieces by weighing

21
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

22
**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

23
Gases and the Mole

24
**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.

25
**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

26
**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)

27
**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

28
**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.

29
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

30
**Quizdom Find the density of CH4 at STP.**

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

31
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

32
**All the things we can change**

33
**Mass Volume Moles 6.02 x 1023 Representative Particles Ions Atoms**

PT Moles Representative Particles 6.02 x 1023 Ions Atoms Count

34
**Percent Composition Like all percents Part x 100 % whole**

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

35
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

36
**Getting it from the formula**

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

37
**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

38
**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

39
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

40
**From percentage to formula**

Empirical Formula From percentage to formula

41
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

42
**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.

43
**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.

44
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

45
**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?

46
**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.

47
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?

48
**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?

49
Example Ibuprofen is % C, 8.80 % H, % O, and has a molar mass of about g/mol. What is its molecular formula?

Similar presentations

OK

Chemical Quantities Chemistry Tracy Bonza Sequoyah High School

Chemical Quantities Chemistry Tracy Bonza Sequoyah High School

© 2018 SlidePlayer.com Inc.

All rights reserved.

Ads by Google

Ppt on waves tides and ocean currents indian Ppt on operating system structure Ppt on entrepreneurship development programmes Ppt on mother teresa life Free download ppt on sources of energy Ppt on any man made disaster Gas plasma display ppt online Ppt on medicinal plants in india Ppt on medical abortion Mp ppt online counselling 2012