1. Chapter 14 Solutions

2. Solutions
Solutions are homogeneous mixtures of two or more pure substances. (remember this from chapter 3)
In a solution, the solute is dispersed uniformly throughout the solvent.

4. Aqueous solution-use water as solvent (recall chapter 7)
Dilute solution- contains a small amount of solute
Concentrated solution- contains large amount of solute

5. How Does a Solution Form:
Chemists use the axiom “like dissolves like”:
Polar substances tend to dissolve in polar solvents.
Nonpolar substances tend to dissolve in nonpolar solvents.

6. How Does a Solution Form: chemical explanation
The intermolecular forces between solute and solvent particles must be strong enough to compete with those between solute particles and those between solvent particles.

7. How Does a Solution Form: chemical explanation
As a solution forms, the solvent pulls solute particles apart and surrounds, or solvates, them.

8. How Does a Solution Form: chemical explanation
The more similar the intermolecular attractions, the more likely one substance is to be soluble in another.

9. How Does a Solution Form: chemical explanation
Glucose (which has hydrogen bonding) is very soluble in water (which has hydrogen bonding), while cyclohexane (which only has dispersion forces) is not.

10. it doesn’t mean the substance dissolved.
Student, Beware!
Just because a substance disappears when it comes in contact with a solvent,
it doesn’t mean the substance dissolved.

11. Student, Beware!
Dissolution is disappearance of a substance into solution
Dissolution is a physical change—you can get back the original solute by evaporating the solvent.
If you can’t, the substance didn’t dissolve, it reacted.

12. Types of Solutions

13. Types of Solutions Saturated
Solvent holds as much solute as is possible at that temperature.
Dissolved solute is in dynamic equilibrium with solid solute particles.

14. Types of Solutions Unsaturated
Less than the maximum amount of solute for that temperature is dissolved in the solvent.

15. Types of Solutions Supersaturated
Solvent holds more solute than is normally possible at that temperature.
These solutions are unstable; crystallization can usually be stimulated by adding a “seed crystal” or scratching the side of the flask.
Example: ornament experiment

16. Solubility Charts

17. WhY do you keep highlighting Temperature?
Generally, the solubility of solid solutes in liquid solvents increases with increasing temperature.

18. WhY do you keep highlighting Temperature?
The opposite is true of gases:
Carbonated soft drinks are more “bubbly” if stored in the refrigerator.
Warm lakes have less O2 dissolved in them than cool lakes.

19. How does the solubility curve relate to saturation?
above the line- supersaturated (holding more than it can) on the line- saturated (can not hold anymore) below the line- unsaturated (can hold more solute)
supersaturated
Solute (g) per g H2O
saturated
unsaturated
temperature

20. Ways of Expressing Concentrations of Solutions

21. Mass Percentage mass of A in solution  100 total mass of solution

22. Parts per Million and Parts per Billion
Parts per Million (ppm)
mass of A in solution
total mass of solution
ppm =
 106
Parts per Billion (ppb)
mass of A in solution
total mass of solution
ppb =
 109

23. Mass Percent sample problem
A solution is prepared by dissolving 1.0g of sodium chloride in 48 g of water. The solution has a mass of 49 g, and there is 1.0g of solute (NaCl) present. Find the mass percent of solute.

24. Sample problem #2
A solution is prepared by mixing 1.00g of ethanol, C2H5OH, with g of water. Calculate the mass percent of ethanol in this solution.

25. Molarity

26. Molarity
molarity is the concentration of a solution given in gram moles of solute per liter of solution.
Ex. A 5M solution of hydrochloric acid has a higher concentration of hydrochloric acid than a 0.2M solution
The higher the molarity in this case the more caution needed.

27. Molarity (M) mol of solute M (mol/L)= L of solution
Because volume is temperature dependent, molarity can change with temperature, but we will save those calculations for AP chemistry

28. Molarity = 0.192 M NaOH 11.5 g NaOH x 1 mol NaOH 40.0 g NaOH
Convert mass of solute to moles (using molar mass of NaOH). Then we can divide by volume
molar mass of solute = 40.0 g
11.5 g NaOH x 1 mol NaOH
40.0 g NaOH
0.288 mol NaOH
1.50 L solution
= mol
NaOH
= M NaOH

29. Molarity
Calculate the molarity of a solution prepared by dissolving 1.56 g of gaseous HCl into enough water to make 26.8 mL of solution.
Given: mass of solute (HCl) = 1.56 g
volume of solution = 26.8 mL

30. Molarity
Calculate the molarity of a solution prepared by dissolving 11.5g of solid NaOH in enough water to make 1.50 L of solution.
Given:
mass of solute = 11.5 g NaOH
vol of solution = 1.50 L
molarity is moles of solute per liters of solution

31. dilution

32. dilution Diluting a solution:
reduces the number of moles of solute per unit volume
the total number of moles of solute in solution does not change

33. M1 = molarity of stock solution (initial)
Diluting solutions
M1V1 = M2V2
M1 = molarity of stock solution (initial)
V1 = volume of stock solution (initial)
M2 = molarity of dilute solution
V2 = volume of dilute solution

34. M1V1 = M2V2
How many milliliters of aqueous 2.00M MgSO4 solution must be diluted with water to prepare mL of aqueous 0.400M MgSO4?
M1 = 2.00M MgSO4
M2 = 0.400M MgSO4
V2 = mL MgSO4
V1 = ?

35. M1V1 = M2V2 Solve for V1 V1 = M2 x V2 M1 0.400M x 100.00 mL 2.00M