1. Unit 5 – Chap 10 & 16 Solutions & Empirical Formulas (last but definitely not least!)
April

2. Solutions and Entropy

4. What do you know already about solutions?
Solutions are:
Solutions of gases, liquids or solids are everywhere
Homogenous mixture
Solute and Solvent
Is this a solution:
AIR
CARBON DIOXIDE GAS
CARBONATED BEVERAGE
SALAD DRESSING
14 KARAT GOLD
24 KARAT GOLD

5. Solutions
A solution is a homogenous mixture of 2 or more substances mixed evenly at a particle level. One part is regarded as the solvent and the others as solute

6. Solution Terminology Solute – substance being dissolved (lesser amt)
Solvent – component whose physical state remain the same (greater amt)
Practice: WHICH IS A SOLUTE AND WHICH IS THE SOLVENT?
2 OUNCES OF OIL & 2 GALLONS OF GASOLINE
CARBON DIOXIDE IN A SELTZER WATER
Dissolve – to mix uniformly and completely at the molecular level
Miscible – liquids that can dissolve in one another form solution
Immiscible - liquids that don’t dissolve in one another
Aqueous solution (aq)

7. Solutions – phet sim
Solubility measures the maximum amount of solute that can be dissolved in a solvent at a given temperature. (measures maximum concentration)

8. Ionic solutes
Called hydrated ions when the solvent is water

9. The dissolving process
Watch NaCl dissolve in water
Watch another version

10. The dissolving process
When an ionic compound dissolves in water:
ionic bonds are broken
attraction between water molecules are broken
attraction between water and ions are formed

11. What happens during dissolving?
How does arrangement of particles change?
How does motion of particles change?

12. Lattice energy and hydration energy

13. Why do solutions form? One major driving force is energy
Is dissolving of an ionic compound endothermic or exothermic? Why?
Example of exothermic dissolving:
dissolving LiCl in water
Example of endothermic dissolving:
dissolving NH4Cl in water
Which bonds/attractions broken during dissolving?
Which bonds/attractions formed during dissolving?
Which are stronger?

14. Entropy: A new concept that explains a lot
Due to the laws of probability, the entropy of the universe is constantly increasing and never decreases.
Processes that would lead to a decrease in entropy of the universe are not possible.
The universe can be divided into the system and its surroundings.

15. Why do things happen?
There are only two reasons for an event to occur (including a chemical process):
Entropy of the system increases
Entropy of the surroundings increases (exothermic reactions are an increase in the surroundings’ entropy)

16. Molarity and Dilutions
Chapter 9
Molarity and Dilutions

17. Solution concentration- terms
Dilute
Small amount of solute
Concentrated
Large amount of solute
Stock - routinely used solutions prepared in concentrated form.

18. How do you quantify concentration of a solution?
Concentration = amount of solute per amount of solution (or solvent)
Molarity = moles solute / liters solution
Example problems:
What is the molarity of a solution if 23 g of KCl are dissolved to make 4 L of solution?
How many grams of KCl should be dissolved to make 500. mL of a 0.10 M solution?

19. Solute concentration expressions
There are a few different ways to calculate the concentration of the solution.
In Biology, we used:
Mass percent: (mass solute / mass of solution) * 100
This year in Chemistry, we will also use:
Molarity(M): moles solute / Liter solution

20. Molarity (M) mol of solute M = L of solution 6 moles of HCl 3 M HCl =
Molarity (M) = moles of solute per volume of solution in liters:
Because volume is temperature dependent, molarity can change with temperature.
mol of solute
L of solution
M =
6 moles of HCl
2 liters of solution
3 M HCl =

21. Concentration: molarity example
If g of KMnO4 is dissolved in enough water to give 250. mL of solution, what is the molarity of KMnO4?
As is almost always the case, the first step is to convert the mass of material to moles.
0.435 g KMnO4 • mol KMnO = mol KMnO4
g KMnO4
Now that the number of moles of substance is known, this can be combined with the volume of solution — which must be in liters — to give the molarity. Because 250. mL is equivalent to L .
Molarity of KMnO4 = mol KMnO4 = M
0.250 L solution

22. Molarity problems
What is the molarity of a solution if 23 g of KCl are dissolved to make 4 L of solution?
How many grams of KCl should be dissolved to make 500. mL of a 0.10 M solution?

23. Molarity problems
3. A truck carrying 22.5 kL of 6.83 M aqueous hydrochloric acid (HCl) used to clean brick has overturned on the highway, and the acid needs to be neutralized. How many moles of HCl have been spilled?

24. moles BEFORE = moles AFTER
Dilution
What does it mean to dilute a solution? What is added?
Solvent is added to lower the concentration.
When a solution is diluted, what is conserved?
The amount of solute remains constant before and after the dilution.
moles BEFORE = moles AFTER
M1V1 = M2V2

25. Common Terms of Solution Concentration
Stock - routinely used solutions prepared in concentrated form. Concentrated - relatively large ratio of solute to solvent. (5.0 M NaCl) Dilute - relatively small ratio of solute to solvent. (0.01 M NaCl): (MV)initial=(MV)Final
Copyright©2000 by Houghton Mifflin Company. All rights reserved.

26. Dilution problems
1. Suppose you have M NaCl stock solution. How do you prepare 250. mL of M NaCl solution ?
Concentration M NaCl

27. Copyright©2000 by Houghton Mifflin Company. All rights reserved.
Figure 4.10: Steps involved in the preparation of a standard aqueous solution.
Copyright©2000 by Houghton Mifflin Company. All rights reserved.

28. Making a solution staring with a solid

29. How to make solutions Starting with a solid:
Does volume of solute = volume of solution?
What type of glassware is best for accurate measurement of a certain volume of liquid AND mixing a solution?
Starting with a concentrated solution:
How does the number of moles of solute change when a solution is diluted?
What type of glassware is best for accurate delivery of a liquid with a certain volume?

30. Dilution problems
What is the molarity of a solution of ammonium chloride prepared by diluting mL of a 3.79 M solution to 2.00 L?
3. To what volume should 1.19 mL of an 8.00 M solution be diluted in order to obtain a final solution that is 1.50 M?

31. Copyright©2000 by Houghton Mifflin Company. All rights reserved.
Figure 4.12: Dilution Procedure (a) A measuring pipet is used to transfer 28.7mL of 17.4 M acetic acid solution to a volumetric flask. (b) Water is added to the flask to the calibration mark. (c) The resulting solution is 1.00 M acetic acid.
Copyright©2000 by Houghton Mifflin Company. All rights reserved.

32. Dilution When a solution is diluted, _____________ are conserved.
Since M = _________, therefore moles = _______
Example problems
What is the molarity of a solution of ammonium chloride prepared by diluting mL of a 3.79 M solution to 2.00 L?
To what volume should 1.19 mL of an 8.00 M solution be diluted in order to obtain a final solution that is 1.50 M?

33. Composition Calculations

34. Describing the composition of compounds
Percent composition (by mass) used to describe what elements a compound is made of, in what proportions.
Experimentally measured
% mass of element =
According to formula OR
According to experimental results

35. Finding experimental percent composition
A sample of iron oxide with a mass of 4.76 g is found to contain 3.70 g of iron. What is its percent composition?
77.7%

36. Calculating theoretical percent composition from formula
What is the percent composition of aluminum oxide?

37. Using percent composition to calculate mass
Calculate the mass of zinc in a 30.00g sample of zinc nitrate.
10.53g

38. 50.0g of SF6
How many grams of sulfur is that?
146.07g/mol
21.96% S

39. Empirical formula
You can use % composition to find out the formula of an unknown compound!
Empirical = determined experimentally
Empirical formula = simplest mole ratio
Examples of empirical formulas:

40. Empirical formula example #1
A sample of an iron oxide is 69.94% Fe and 30.06% O. What is the empirical formula?
Fe2O3

41. Empirical formula example #2
In lab, a student analyzes a sample of nickel chloride and finds that it contains 35% Ni and 65% Cl. What is the empirical formula?
NiCl3

42. Empirical formula example #3
A piece of copper ore with a mass of grams is found to contain g of copper. The remainder of the mass is fluorine. What is the empirical formula of this ore?
CuF2

43. Practice problems Try these problems in small groups:
What is the percent composition of tin(IV) oxide?
A sample of silver(I) sulfide has a mass of 62.4 g. What mass of each element could be obtained by decomposing this sample?
A sample of copper chloride is found to contain 5.46 g of copper and 6.10 g of chloride. What is its percent composition?

44. Lab procedures
Iodine vapor is toxic – keep the beakers covered and don’t breathe fumes.
Decant:
Gentle heating:
Mass of product:

45. PERCENT YIELD
Theoretical Yield – the “calculated” amount. It is the maximum amount that can be produced by the given reaction.
Reasons for less than 100% yield:
Incomplete rxn of the limiting reagent
Less than ideal reacting conditions
Reversible rxns
Formation of unwanted reaction products
Loss of product in transferring from one vessel to another
Actual Yield (experimental) - the quantity of the product that is actually obtained from the reaction.