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**Please pick up the materials in the front and take a seat**

Happy New Year! Please pick up the materials in the front and take a seat

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How was my break?

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**Policies and Procedures**

- Test Corrections - Pink Sheets & Late Work - Notes & Practice Problems

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Chapter 8 BEAT Sheet Calculate the percent composition of any element in a compound Determine the empirical formula of a compound given its percent composition Determine the molecular formula of a compound given its percent composition and molar mass Determine the empirical formula of a hydrate from lab data Perform mole road calculations to determine atoms, molecules, ions, mass, or moles from relevant data Describe the labs you performed this unit and what the key concept/learning was from each lab

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**We have talked about this**

What is a mole? We have talked about this 6.02 X 1023 times!

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What is a mole? A mole is the base unit used to measure the amount of a substance, abbreviated mol. One mol of anything is equal to x 1023 pieces of that thing.

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Mole Practice 1. How many moles in 5.4 x 1020 atoms of Pb?

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Mole Practice 2. How many atoms in 3.5 moles of Ag?

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Mole Practice 3. How many total molecules in 1 mole of glucose: C6H12O6?

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**Molecular and Formula Weight**

Last semester we learned about molecular weight and touched on formula weight. It is important to remember that a single atom’s weight in amu is the same as one mole of that atom in grams. Example: The weight of a single atom of aluminum is grams. Therefore 1 mol of Al weighs grams.

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**Molecular and Formula Weight**

This is known as the atomic mass: the mass of 1 mole of atoms of any element. This is the number that’s on the periodic table. But atoms are rarely by themselves—they are usually in compounds. How do you find the mass of a mole of a compound??

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**Molecular and Formula Weight**

Example: Find the molecular mass of CO2 C: 1 x = g O: 2 x = g g

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**Molecular and Formula Weight**

Example: Find the formula mass of (NH4)CO3 N: 1 x = g H: 4 x = g C: 1 x = g O: 3 x = g g

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**Molecular and Formula Weight Practice Problems**

4. Find the molecular mass of C6H12O6

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**Molecular and Formula Weight Practice Problems**

5. Find the molecular mass of dinitrogen pentaoxide

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**Molecular and Formula Weight Practice Problems**

6. Find the Formula mass of Na2SO4

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**Molecular and Formula Weight Practice Problems**

7. Find the Formula mass of potassium phosphate

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**Molar Mass of a Substance**

When trying to find the mass of one mole of a substance we will make life easier by just referring to the molar mass of a material. It can be an atom, a molecule, or an ionic compound. We use the same technique to calculate the mass of a mole of O2, Pb, CH4, or NaCl.

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**Molar Mass Practice Problems**

8. Find the molar mass of O2

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**Molar Mass Practice Problems**

9. Find the molar mass of Pb

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**Molar Mass Practice Problems**

10. Find the molar mass of CH4

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**Molar Mass Practice Problems**

11. Find the molar mass of NaCl

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**Molar Mass Practice Problems**

12. When do you use atomic mass, molecular weight, and formula weight to find the molar mass?

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Molar Mass Conversion Since molar mass gives you the amount of grams in one mole of a substance you can use the molar mass to convert from grams to moles or vice versa.

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** 56.8 𝑔 𝑀𝑛𝑂2 1 1 𝑚𝑜𝑙 𝑀𝑛𝑂2 86.9368 𝑔 𝑀𝑛𝑂2 = .653 𝑚𝑜𝑙 𝑀𝑛𝑂2**

Molar Mass Conversion Examples: How many moles in 56.8 g of MnO2? 𝑔 𝑀𝑛𝑂 𝑚𝑜𝑙 𝑀𝑛𝑂 𝑔 𝑀𝑛𝑂2 = .653 𝑚𝑜𝑙 𝑀𝑛𝑂2

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**Molar Mass Conversion Examples: How many grams in 3.4 moles of C6H6? **

3.4 𝑚𝑜𝑙 𝐶6𝐻 𝑔 𝐶6𝐻6 1 𝑚𝑜𝑙 𝐶6𝐻6 =260 𝑔 𝐶6𝐻6

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**Molar Mass Conversion Practice Problems**

13. How many moles are in g of HNO3

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**Molar Mass Conversion Practice Problems**

14. How many moles are in 19.5 grams of NaCl

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**Molar Mass Conversion Practice Problems**

15. How many grams are in 15 mols of AlPO4

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**Molar Mass Conversion Practice Problems**

16. How many grams are in 5.34 mols of KBr

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**The Volume of a Mole of Gas**

Until this point we have only converted from mass to moles, but what happens when we convert from volume to moles? Describe how we can measure the mass of a solid substance: Describe how we can measure the mass of a liquid substance: Can we measure the mass of a gaseous substance?

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**The Volume of a Mole of Gas**

When we convert from volume to moles or moles to volume we use a constant of 22.4 liters per mole. This constant tells us that one mole of a gas has a volume of 22.4 liters. When using this constant we assume the gas is at STP (standard temperature and pressure) which is 1 atm and 0C

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**The Volume of a Mole of Gas**

Also remember that if there is one mole of a sample of gas, that sample must have 6.02 x 1023 atoms in it! But since gases have different masses than their solid form, one mole of a gas does not weigh the same as its solid state!

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**The Volume of a Mole of Gas**

Example: If the density of a gas is 2.3 g/L, what is its molar mass? 2.3 𝑔 1 𝐿 𝐿 1 𝑚𝑜𝑙 =51.52 𝑔/𝑚𝑜𝑙

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**The Volume of a Mole of Gas**

Example: What is the density of F2 at STP?

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**The Volume of a Mole of Gas Practice Problems**

17. Since we know that the gaseous state of a substance does not have the same mass as the solid state of the same substance, what can we say about the density of the two?

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**The Volume of a Mole of Gas Practice Problems**

18. How many moles are in 22.4 Liters of O(gas)?

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**The Volume of a Mole of Gas Practice Problems**

19. How many milliliters of Neon gas are in 2 mols?

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**The Volume of a Mole of Gas Practice Problems**

20. What is the molar mass of a gas that has a density of g/L? Which diatomic gas do you think it is?

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**The Volume of a Mole of Gas Practice Problems**

21. What is the density of Cl2 at STP?

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22.4 g/L Molar Mass Volume Avogadro’s number Moles Atoms

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**Mole Conversion Practice Problems**

22. How many moles in 9.0 liters of a gas at STP?

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**Mole Conversion Practice Problems**

23. How many moles in 15 g of NaOH?

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**Mole Conversion Practice Problems**

24. How many liters would 2.9 moles of CO2 occupy at STP?

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**Mole Conversion Practice Problems**

25. How many grams would 30 L of CO2 gas weigh?

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**Mole Conversion Practice Problems**

26. How many atoms are in 50 L of H2 gas?

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**Calculating Percent Composition**

Percent Composition is the percent by mass of each element in a compound. You find the percent composition by dividing the mass of the element by the total mass of the molecule. If you think back to the separation lab we have done this already by calculating the % of salt, % of sand, and % of iron!!

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**Calculating Percent Composition**

Example: How many students are in this class? Find the percentage of Males and the percentage of Females.

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**Calculating Percent Composition**

Example: What percentage of students in the class are 15? 16? 17?

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**Calculating Percent Composition**

Example: What is the percent composition of each element in C6H12O6?

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**Calculating Percent Composition**

Example: What is the percent by mass of each element in H2O?

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**Calculating Percent Composition**

Example: How many grams of nitrogen are in 59 g of AgNO3?

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**Percent Composition Practice Problems**

27. Which compound has the smallest percentage of Chlorine? HCl KCl LiCl NaCl

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**Percent Composition Practice Problems**

28. Of the following compounds, which does Strontium have the highest percent composition? SrCl2 SrI2 SrO SrS

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**Percent Composition Practice Problems**

29. What is the percent composition of Oxygen in the compound propanal (CH3CH2CHO)?

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**Percent Composition Practice Problems**

30. What is the percent by mass of Oxygen in H2SO4?

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**Percent Composition Practice Problems**

31. What is the total mass of oxygen in 1.00 mol of Al2(CrO4)3?

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**Percent Composition Practice Problems**

32. What is the percent composition of Carbon in CO?

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**Calculation Empirical Formulas**

An empirical formula is the lowest whole number ratio of the atoms of the elements in a compound. The empirical formula may or may not be the same as the molecular formula.

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**Calculation Empirical Formulas**

Example: The molecular formula for CO2 is the same as the empirical formula for CO2 because the atoms are in the lowest whole number ratio.

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**Calculation Empirical Formulas**

Example: The molecular formula for C2H2, could be that, or it could be C6H6. But the empirical formula for both is CH, since that is the lowest whole number ratio.

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**Calculation Empirical Formulas**

It is important to understand that the empirical formula gives you an idea of the ratio of between elements within a compound.

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**Calculation Empirical Formulas**

Example: Calculate the empirical formula of a compound made up of: 32.00% C 42.66% O 18.67 % N 6.67% H

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**Calculation Empirical Formulas Practice Problems**

33. A compound contains 40% calcium, 12% carbon, 48% oxygen by mass. What is the empirical formula for this compound?

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**Calculation Empirical Formulas Practice Problems**

34. Write the empirical formula for the following compound: P4O10

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**Calculation Empirical Formulas Practice Problems**

35. A compound is 86% carbon and 14% hydrogen by mass. What is the empirical formula for the compound?

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**Calculation Empirical Formulas Practice Problems**

36. Which of the following is an empirical compound? C4H10 C6H12O6 P2O5 C2H2O2

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**Calculation Empirical Formulas Practice Problems**

37. A compound was analyzed and was found to contain 75% carbon and 25% hydrogen by mass. What is the compounds empirical formula?

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**Calculation Empirical Formulas Practice Problems**

38. What is the simplest ration of nitrogen to oxygen in the compound nitrogen (IV) oxide

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**Calculation Molecular Formulas**

Knowing that the Molecular formula of a substance is not always equal to the Empirical Formula of the substance, what can we say about Molecular Weight vs Empirical Formula Weight?

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**Calculation Molecular Formulas**

We can determine the molecular formula we first know the molar mass and the empirical formula of the substance. We then compare the molar mass to the Empirical Formula Weight to determine the ratio.

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**Calculation Molecular Formulas**

Examples: The molar mass of a substance is 60g and its empirical formula is CH4N. Find the Molecular Formula.

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**Calculation Molecular Formulas**

Examples: The molar mass of a substance is 78g and its empirical formula is CH. Find the Molecular Formula.

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**Calculation Molecular Formulas**

Examples: Find the molecular formula for a compound that has: 54.5% C 13.6% H 31.8 % N molar mass = 88g

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**Calculation Molecular Formulas Practice Problems**

39. What’s the molecular formula of a compound given the empirical formula P2O5 and molar mass 284 g?

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**Calculation Molecular Formulas Practice Problems**

41. What is the molecular formula of a compound that has a molecular mass of 42 g and empirical formula of CH2?

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**Calculation Molecular Formulas Practice Problems**

42. Find the molecular formula for a compound that has: 54.5% C 13.6% O 31.8 % N molar mass = 176 g/mol

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**Calculation Molecular Formulas Practice Problems**

43. Find the molecular formula for a compound that has: 25.9% N 74.1% H molar mass = 108 g/mol

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**Calculation Molecular Formulas Practice Problems**

44. Find the molecular formula for a compound that has: 22.8 g Ba 2.00 g C 8.00 g O molar mass = g/mol

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Hydrates Hydrated compounds have water molecules attached to them. CuSO45H2O is copper (II) sulfate pentahydrate. For every formula unit of CuSO4, there are 5 water molecules attached to it. There are always numerous waters for one salt.

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Hydrates Example: Start with 1.54 g wet magnesium phosphate. Dry it so that there are only 1.28 g residue left. Find the formula of the hydrate.

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Hydrates Example: Calculate the percent composition in NaCl ∙ 5H2O?

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**Hydrates Practice Problems**

45. Calculate the percent composition in BaCl2 ∙ 2H2O

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**Hydrates Practice Problems**

46. What is the percent by mass of the water in the hydrate Na2CO3 ∙ 10H2O?

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**Hydrates Practice Problems**

47. If you start with 1.54 g wet magnesium phosphate. Dry it so that there are only 1.02 g residue left. Find the formula of the hydrate.

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**Hydrates Practice Problems**

48. Start with 1.62 g of hydrated cobalt (II) chloride. After heating, there are 0.88 g of solid reside. Find the formula of the hydrate.

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Chapter 10 Chemical Quantities

Chapter 10 Chemical Quantities

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