1. Solutions & Acids and Bases
Units 15 & 16
Solutions & Acids and Bases

2. Solutions All solutions are composed of two parts:
The solute and the solvent.
The substance that gets dissolved is the solute
The substance that does the dissolving is the solvent (Usually present in the larger amount)
***A solution may exist as a solid, liquid or gas depending on the state of the solvent.

3. Types of Solutions Gas in liquid Solid in solid
Example – soda water
Solute: carbon dioxide (gas) Solvent: water (liquid)
Solid in solid
Example – Steel
Solute: carbon
Solvent: iron

4. Types of Solutions Gas in gas Example - Air Solute: Oxygen
Solvent: Nitrogen

5. Types of Solutions Liquid in liquid Example – Vinegar Solid in liquid
Solute: Acetic acid
Solvent: Water
Solid in liquid
Example – Ocean Water
Solute: Sodium Chloride (solid)
Solvent: Water (liquid)

6. Aqueous Solution Any mixture where water is the solvent.
Something is dissolved in water

7. Solubility Soluble - a substance that dissolves in another substance.
Insoluble - a substance that does not dissolve in another substance.

8. Solubility Immiscible - two liquids that are insoluble in each other.
Miscible - two liquids that are soluble in each other.

9. Solvation (Hydration)
Solvation – a process that occurs when an ionic solute dissolves in a solvent
Solvation Video

10. “Like Dissolves Like” Aqueous solutions of ionic compounds:
Solvents of a specific polarity or type will dissolve solute of similar polarities or types!
Aqueous solutions of ionic compounds:
The charged ends of the water molecules attract the positive and negative ions making up an ionic solid, forcing them to separate.
Aqueous solutions of molecular compounds:
Molecular compounds that have polar sections easily form aqueous solutions with water.

11. Solubility
Solubility – the maximum amount of solute that will dissolve in a given quantity of solvent at a specific temperature and pressure to produce a saturated solution
Units for solubility: grams of solute per 100 g solvent
Example: At 20˚C, NaNO3 has a solubility of 74 g/100 g H2O

12. Solubility
Saturated Solution - contains the maximum amount of dissolved solute
Unsaturated Solution - contains less than the maximum amount of dissolved solute
Supersaturated Solution – contains more solute than can theoretically be dissolved at a given temperature

13. Supersaturated Solutions
How can you dissolve more solute than possible??

14. Solubility Curves
Solubility of a solid generally increases with increasing temperature
The higher the temperature, the greater amount of solute that will be dissolved in the solvent
Solubility can be represented in a chart called a solubility curve

15. Solubility Curves

16. Factors that affect solubility
Temperature-
Generally, as temperature increases, more solid solute will dissolve in the same amount of liquid solvent. The opposite is true for gaseous solutes.

17. concentration
Concentration –the measure of the amount of solute dissolved in a given quantity of solvent

18. Solution Concentration
Expressing concentration:
Concentration Description
Ratio
Percent by mass
Percent by volume
Molarity

19. Percent By Mass
Percent by mass -ratio of the solute’s mass to the solution’s mass expressed as a percent.
Example: An aquarium contains 3.6 g NaCl per g of water. What is the percent by mass of NaCl in the solution?

20. Percent By Volume
Percent by volume-ratio of the volume of the solute to the volume of the solution expressed as a percent.
Example: What is the percent by volume of ethanol in a solution that contains 35 mL of ethanol dissolved in 115 mL of water?

21. molarity Molarity (M) = moles of solute liters of solution
Moles of solute dissolved in 1 liter of solution
Example: 0.23 M solution = 0.23 moles of solute dissolved in 1 L of solution

22. molarity M is read as “molar” when next to a number
4 M HCl = 4 molar hydrochloric acid
Keep in mind that the liters of solution takes into account the volume of the solute AND the volume of the solvent

23. Example: What is the molarity of a solution that contains 0
Example: What is the molarity of a solution that contains 0.65 mol of CuCl2 in 500 mL of water?

24. Example: What is the molarity of a solution that contains 5
Example: What is the molarity of a solution that contains 5.10 g of glucose (C6H12O6) in mL of solution?

25. Preparing Molar Solutions
How is a solution of known molarity made?
Convert moles of solute to grams and measure the amount out.
Add solvent so that the total volume of the solution is 1 L.
For any volume other than 1 L we must adjust the amount of solute needed by multiplying it by the fraction of a liter of solution we need.

26. Example
Example: How many grams of CaCl2 would be dissolved in 1.0 L of water to make a 0.10 M solution of CaCl2?

27. Example
Example: How many grams of NaOH are in 250 mL of a 3.0 M NaOH solution?

28. dilutions
Dilutions are used to decrease the concentration (or molarity) of a solution
M1V1=M2V2

29. dilutions Steps to Performing a Dilution
Calculate how many mL of the original (stock) solution to start with
Measure out the volume of stock solution (using a graduated cylinder or a pipet) and place in appropriately sized volumetric flask
Add water to the mark on flask

30. dilutions

31. Example:
What volume, in milliliters, of 2.00 M CaCl2 is needed to make 0.50 L of M CaCl2 solution?

32. Colligative Properties
Colligative means “depending on the collection.”
Depends only on the number of dissolved particles, not on the identity of dissolved particles.
Includes vapor pressure lowering, boiling point elevation, freezing point depression, and osmotic pressure

33. Electrolytes and Colligative Properties
Electrolyte: Soluble ionic compounds. When they dissolve in solution, they dissociate into their component ions and conduct electricity.
Ex: NaCl (s)  Na+ (aq) + Cl- (aq)
Covalent molecules in aqueous solution:
Covalent particles do not dissociate when in solution, so the # of molecules = the # of particles.

34. Examples of Colligative Properties
Boiling Point Elevation
Boiling occurs when vapor pressure equals atmospheric pressure.
Boiling point of a solution is higher than the boiling point of the pure solvent.
Dissolving substances increases the boiling point of a solvent.

35. Freezing Point Depression
Freezing point of a solution is lower than the freezing point of the pure solvent.
Dissolving substances lowers the freezing point of a solvent.
Ex: Icy pavement - throw down CaCl2 or NaCl, and the water will then freeze at a lower temperature

36. Ex: Antifreeze: a solution of ethylene glycol in water
1. Prevents car’s radiator from freezing in the winter.
2. Prevents car’s radiator from boiling over in the summer
The more ethylene glycol in the water, the lower the freezing point, and the higher the boiling point.

37. Acids & Bases

38. Properties of Acids Physical: Chemical: Litmus Indicator:
Taste sour
Chemical:
React with metals to produce H2 gas
Neutralized when reacted with a base
Litmus Indicator:
Turns blue litmus paper red
Ions in Solution:
H+, H3O+ (hydronium ion)

39. Properties of Bases Physical: Chemical: Litmus Indicator:
Taste Bitter
Slippery
Chemical:
Neutralized when reacted with an acid
Do not react with metals
Why are bases used as drain cleaners?
Litmus Indicator:
Turns red litmus paper blue
Ions in Solution:
OH-1

40. Arrhenius Acids & Bases
Arrhenius Model:
ACIDS: Acids contain the H+ ion
Ex.) HCl, HBr, HNO3
BASES: Bases contain the OH-1 ion
Ex.) NaOH, KOH, Ca(OH)2

41. Bronsted-Lowry Acids & Bases
Bronsted-Lowry Model:
For every acid, there must be a base
Acid = proton donor
Base = proton acceptor
HCl (aq) + NH3 (aq)  NH4+ (aq) + Cl-1 (aq)

42. Conjugate Pairs
NH3 / NH4+ is a conjugate pair — related by the gain or loss of H+
Every acid has a conjugate base, formed when H+ is removed from the acid.
Every base has a conjugate acid, formed when H+ is added to the base.

43. Conjugate Pairs Identify the conjugate acid-base pairs: HCl in water
NaOH in water
NH3 in water

44. Identify the following as monoprotic or polyprotic:
Types of Acids
Monoprotic and Polyprotic Acids
Acids can contain 1 or more hydrogens that are acidic
**Not ALL hydrogens are acidic (Ex. Vinegar)
Identify the following as monoprotic or polyprotic:
HNO3, H2SO4, HClO, HClO4, H3PO4, HC2H3O2

45. Strength of Acids/Bases
Strengths of Acids
Strong Acid Give off LOTS of H+
100% Dissociation
Strong Acids: HCl, HI, HBr, HNO3, H2SO4, HClO4
That’s it! Everything else is “weak”
Weak Acid Give off smaller amounts of H+
Equilibrium occurs (breaks apart and then recombines)
Not all H+ ions separate (not 100% dissociation)

46. Strength of Acids/Bases
Strengths of Bases
Strong Base Give off LOTS of OH-1
100% Dissociation
Generally, Group I, II Hydroxides (except H, Be, Mg)
Ex.) Ca(OH)2, NaOH
Everything else is “weak”
Weak Base Give off smaller amounts of OH-1
Equilibrium occurs (breaks apart and then recombines)
Not 100% dissociation

47. Strength of Acids/Bases
Strong or weak vs. concentrated and dilute
Strong/weak tells you how much it dissociates
Concentrated/dilute indicates the concentration (amount of solute in the solvent)

48. pH pH & pOH pH Scale pH tells us the acidity or basicity of a solution
Based on measuring the [H+] (a.k.a. [H3O+])
pH Scale
Ranges 0 to 14
Acid ~ 0 to 7
Bases ~ 7 to 14

49. Definition: Hydronium Ion
In aqueous solution, H+ does NOT exist!
Note: In problems, [H+] = [H3O+]
H+ + H2O  H3O+
(hydronium ion)

50. pH pH = - log [H3O+] pOH = - log [OH-] pH + pOH = 14
Make sure you have the negative sign!
Find the “log” function on your calculator!
pH + pOH = 14
[H+][OH-] = 1.0 x 10-14

52. pH Calculations
What is the pH and pOH for a solution with a H+ concentration of 3.0 x 10-6 M H+?

53. pH Calculations
What is the H+ and OH- concentration of blood with a pH of 7.40?

54. Neutralization Reactions
Reaction in which acid and base react to neutralize one another
Acid + Base  Water + Salt
***Salt = Any ionic compound formed as a
by-product of an acid-base reaction

55. Neutralization Acid-base Titration: Definition:
Lab technique which allows you to get moles of acid and base EXACTLY equal to another
Complete neutralization
Allows you to calculate the concentration of an unknown acid or base

56. Definitions
The titrant is the substance of known concentration used to determine the unknown concentration of the other substance.
An indicator- substance that changes color at a certain pH—is added to tell us when the neutralization is complete.
Example: Phenolphthalein undergoes a color change between pH 8 and 10
clear in acid
Light pink in neutral
Dark pink in base

57. Neutralization Procedure:
Add known volume of acid or base to Erlenmeyer flask
Add a known concentration of the other to a buret
Add an indicator to the flask
Slowly dispense titrant (what you’re adding with a buret) into the flask
Stop when 1 drop of titrant causes the indicator to switch from one color to another

58. Neutralization Equivalence point: pH at which amount of acid = base
Indicator: Compound that changes color due to a change in pH
Common Indicators and pH Range
Litmus: 5.5 to 8.0 (red= acid, blue = base)
Phenolphthalein: 8.2 to 10.6 (colorless to magenta)
End point: Point at which the volume of titrant added makes the amount of acid and base are equal and the indicator changes color