1. UNIT 12: ACIDS, BASES, and SALTS

2. Review!
What is an electrolyte?
What are examples?

3. Aim # 1 How do we define acids and bases?
Acidity: measure of the hydrogen (H+), hydronium (H3O+)ion concentration of a solution. Dissolve in solution to produce ions.
Common acids are found on Table K
Properties of ACIDS:
In aqueous solution they are electrolytes, conduct electricity
Strong acids: good conductors, 100% dissociation
Weak acids: poor conductors, no dissociation
2. Sour Taste examples:
3. React with bases (hydroxides) to produce water and salt
Neutralization reaction:

4. 4. React with certain metals to produce H2 gas (table J)
5. Cause acid-base indicators to change color (Table M)
an indicator is:
litmus paper:
phenolphthalein:
6. pH less than 7
7. Acids can be formed by a reaction of gaseous oxides with water
8. General formula is HX (where X represents an anion such as Cl-

5. In aqueous solutions they are electrolytes, conduct electricity
Alkalinity is a measure of the hydroxide ion concentration of a solution, dissolve in solution to produce ions
Common bases are found on Table L
Properties of BASES
In aqueous solutions they are electrolytes, conduct electricity
a. strong bases:
b. weak bases:
Bitter taste examples:
Slippery/soapy feel examples:
React with acids to form salt and water

6. 5. Cause acid-base indicators to change colors
litmus:
phenolphthalein
6. pH greater than 7
7. Formed when Group 1 and 2 metals react with H2O, H is released too
8. General formula is XOH (where X represents a cation such as Na+)

7. AIM# 2: What are the acid and bases theories?
1. Arrhenius Theory
an acid is defined as a substance whose water (aqueous) solution contains or yields hydrogen or hydronium ions (H+, H3O+) as the only positive ion in solution
Example H2SO4  H+ + H+ + SO4-2
HCl  H Cl-
**Strong Acids will dissociate completely in water**
Examples:

8. **Remember not all OH groups are bases, some are alcohols
b) A Base is defined as a substance whose water (aqueous) solution contains or yields OH- as the only negative ion. The presence of the OH- group makes the base an electrolyte.
**Remember not all OH groups are bases, some are alcohols
CH3OH (methanol) or similar compounds are not bases!
CH3COOH is an acid! Check table K

9. practice Which substance is classified as an Arrhenius acid?
a. HCl b. NaCl c. LiOH d. KOH
According to Table J which metal with react with 0.1 M HCl?
a. Au b. Ag c. Hg d. Mg
Which solution turns phenolphthalein pink and red litmus blue?
a. HCl b. CO2 c. NaOH d. CH3OH
Which ion is produced when an Arrhenius base is dissolved in water?
a. H+ is the only positive ion in solution c. OH- is the only negative ion in solution
b. H3O+ is the only positive ion in solution d. H- is the only negative ion in solution
Which species can conduct an electric current?
a. NaOH(s) b. H2O(s) c. CH3OH(aq) d. HCl(aq)

10. 2. Alternate Acid- Base Theory (Bronsted – Lowry; BAAD
This theory explains the behavior of WEAK acids and WEAK bases
Acids are proton (H+) donors
The acid will donate an H + to the base
Bases are proton (H+) acceptors
The base will accept the H+ ion from the acid

11. CONJUGATE ACID-BASE PAIRS:
A conjugate pair refers to acids and bases with common features. These common features are the equal loss/gain of protons between the pairs. Conjugate acids and conjugate bases are characterized as the acids and bases that lose or gain protons.  In an acid-base reaction, an acid plus a base reacts to form a conjugate base plus a conjugate acid
Acid + Base→ Conjugate Base + Conjugate Acid

12. Note: the acid always has one more H than the base!
The conjugate acid of a base is formed when the base gains a proton. Refer to the following equation:

13. Amphoteric Substances:
Examples:
Citric acid, fruits, sour candies
Strong acid (Strong electrolyte) HI, HCl, HNO3 ,H2S04 weak acid (weak electrolyte) H2CO3 (soda), HC2H3O2 (found in vinegar)
7. Acid rain, burning of fossil fuels releases nonmetallic oxides into atmosphere, combine with H2O form weak acids
*usefulness of acid/base indicators indicate or tell you what type of solution you are dealing with

14. Practice : HCl + H2O  H3O+ + Cl- NH3 +H2O  NH4++ OH- HC2H3O2 + H2O  H3O+ + C2H3O2-

15. AIM #3 : How can we determine the strength of acids and bases?
pH scale: direct measurement of H+ ion concentration in a solution.
The pH scale indicates the strength of an acid or base
POWER OF HYDROGEN: developed to express [H+] as # from 0-14
0= strongly acidic (0-6 acidic)
7= neutral
14= strongly basic (8-14 basic)
Different acids because they have different strengths
Ex) if 100 molecules of HCl dissolved in water all 100 molecules ionize and form ions
Ex) other weaker acids, bases maybe 1 out of 100 molecules will ionize and form ions

16.  LOGARITHMIC SCALE
Each change of a pH unit signifies a tenfold change in the concentration of the hydrogen ion, using the power of 10.
Ex) Concentration of [H+] is 10x greater in a solution with a pH = 5 than pH = 6
105 is ten times stronger then 106
A pH of ________ has 10x as much H+ in a solution than a pH of ______________
A pH of 1 is ________________ stronger than a pH of 3
A pH of 2 is _________________ stronger than a pH of 5
When the pH increases by 1 unit (from 5-6) the H+ concentration __________________ by 10.
A pH of 6 is ______________ weaker than a pH of 5.

17. Different acids because they have different strengths
RELATIVE STRENGTHS OF ACIDS AND BASES
Strength is determined by how much it will ionize/dissociate. Highly ionized acids and bases will produce large numbers of ions
Different acids because they have different strengths
Ex) if 100 molecules of HCl dissolved in water all 100 molecules ionize and form ions
Ex) other weaker acids, bases maybe 1 out of 100 molecules will ionize and form ions

18. Boric Acid – eye washing
Strong acids
Strong Acids (H+ or H3O+)
Found at the top of Table K
Low pH number (1, 2 or 3)
The more oxygen present in the polyatomic ion of an oxyacid (acid containing oxygen), the STRONGER its acid within that group is
HClO, HClO2, HClO3
HCl – dangerous acid
Citric Acid – in fruits
Boric Acid – eye washing

19.  Strong Bases (OH-)
Found at the top of Table L
High pH number (12, 13 or 14)
Hydroxides or oxides of group 1 and 3 metals (except Mg and Be)
Those that are very soluble are very strong

20. Diprotic acid:
Triprotic Acid:

21. Aim# 4 How can we determine if a metal will react with an acid spontaneously?
According to Table J, any metal located above H2 will react with an acid to produce H2 gas and a salt Which metal, Mg or Cu will react with HCl? General Reaction: Metal + Acid  H2 + salt Zn + 2HCl  ____________ + ____________ Ca + 2HCl  ___________ + ___________ Cu + H3PO4  __________ + _____________ Which metals do not react with H2??

22. AIM # 5: How can we determine acidity or alkalinity?
Acid- Base Indicators (Table M) An indicator is a substance that changes color as a result of a pH change. Indicators are chosen based on their range for color change. - Chemicals that have certain coloring depending on pH - pH paper to compare results to chart TABLE M A solution yields the following results when tested with various indicators
1 – calibrate it before use by putting it into two solutions with different pH values

23. Practice: At a pH of 2, determine the color of the following:
Methyl orange: ____ Bromthymol Blue: ____ Phenolphthalein:___
Litmus: _________ Bromcrescol Green: ___ Thymol Blue: ____

24. AIM# 7: HOW CAN WE DETERMINE THE CONCENTRATION OF AN ACID OR BASE?
TITRATION:
Is the controlled process of acid-base neutralization used to determine the concentration of an acid or a base
The acid of an unknown molarity is reacted with a carefully measured amount of a base of known molarity to the point of neutralization

25. SOLVING TITRATION PROBLEMS
Endpoint: is the pH of the solution at the desired color (change)
Equivalence Point: the point at which the titrated solution has a pH of 7, when the titration is complete
[OH-] = [H+] – can be detected using probes
Table T Formula: Macid x Vacid = Mbase x Vbase
Use this formula when you are dealing with a titration or neutralization word problem
SOLVING TITRATION PROBLEMS

26. 1. What is the concentration of a solution of HI if 0
1. What is the concentration of a solution of HI if 0.3L is neutralized by 0.6L of 0.2M solution of KOH?
2. What is the concentration of a hydrochloric acid solution if 50.0mL of a M KOH solution are needed to neutralize 20.0mL of the HCl solution of unknown concentration?
3. A particular acid has an H + concentration of 0.1M and a volume of 100mL. What volume of a base with a 0.5M [OH-] will be required to neutralize the reaction?

27. 4. If it takes 15. 0mL of 0. 40M NaOH to neutralize 5
4. If it takes 15.0mL of 0.40M NaOH to neutralize 5.0mL of HCl, what is the molar concentration of the HCl solution?
5. If it takes 10.0mL of 2.0M H2SO4 to neutralize 30.0mL of KOH, what is the molar concentration of KOH?
6. How many mL of 2.0M H2SO4 are required to neutralize 30.0mL of 1.0M NaOH solution?
7. How many mL of 0.10M Ca(OH)2 are required to neutralize 25.0mL of 0.50M HNO3 solution?

28.  Neutralization
This occurs when an acid and base react to form a salt and water
example: antacid for upset stomachs neutralize the acid in the stomach and produces a neutral salt to provide relief
Complete the following reactions
HBr + KOH 
NaOH + HC2H3O2 
KOH + H3PO4 
2HNO3 + Ca(OH)2 

29. PAGE 26 REVIEW OF NAMING ACIDS AND BASES
1. Binary Acids: two elements begin with hydro- and the non-metal will end with –ic, add acid at the end.
(TABLE K)
2. TERNARY ACIDS: cover up the “H” and look at the polyatomic ion. Change the endings of the compounds, and add acid at the end.
-ate to –ic
-ite to –ous
3. BASES: Name metal, add (OH) hydroxide