1. Acids and Bases; pH

2. Acid Base Reactions
We learned that an acid base reaction was a special type of a double displacement reaction.
An acid reacts with a base to produce a salt and water.

3. What is an acid?
According to the Arrhenius acid-base concept, a substance is classified as an acid if it ionizes to form hydrogen(+) ions in aqueous solution. For example, hydrochloric acid reacts with water to form hydrogen ions which are transferred to a water molecule to form a hydronium ion (H3O+).

4. What is an acid?
Other systems classify substances as acids if they act as proton donors (Bronsted-Lowry theory) or as electron-pair acceptors (Lewis theory).
These two classification methods are not limited to solutions in water, but are used for other solvents as well.

5. Common Acids
Common acids are citric acid (from certain fruits and veggies, notably citrus fruits)
ascorbic acid (vitamin C, as from certain fruits)
vinegar (5% acetic acid)
carbonic acid (for carbonation of soft drinks)
lactic acid (in buttermilk)

6. What is a Base?
Depending on the classification system, bases produce OH- ions in aqueous solutions (Arrhenius acid-base concept), are proton acceptors (Bronsted-Lowry theory), or are electron-pair donors (Lewis theory).
In other words, bases behave oppositely from acids.

7. Common Bases Detergents Soap Lye (NaOH) Household ammonia (aqueous)
Baking soda

8. Acid Strength
The strength of an acid refers to its ability or tendency to lose a proton (H+).
A STRONG acid completely ionizes (dissociates) in a solution to produce hydronium ions (H3O+). It ionizes 100%.
A mole of a strong acid HA dissolves in water to yield one mole of H3O+ and one mole of the conjugate base, A-. Effectively there is no ionized HA remaining. (This reversible reaction has achieved chemical equilibrium.)
HA + H2O ↔ H3O+ + A-

9. Strong Acids These six acids are the strong acids:
hydrochloric acid (HCl)
hydroiodic acid (HI)
hydrobromic acid (HBr)
perchloric acid (HClO4)
nitric acid (HNO3)
sulfuric acid (H2SO4).

10. CH3COOH + H2O ↔ CH3COO- + H3O+
Weak Acids
Weak acids do not ionize fully when dissolved in water. An example is acetic (or ethanoic) acid.
CH3COOH + H2O ↔ CH3COO- + H3O+
At any one time, only about 1% of the acetic acid molecules have converted into ions. The rest remain as simple acetic acid molecules in solution.
Hydrogen fluoride dissolves in water to produce hydrofluoric acid (HF), a weak inorganic acid.
Formic acid (HCOOH) is a weak organic acid (think ants!).
Most organic acids are weak (e.g., acetic, formic, citric, lactic, carbonic).

11. Strong bases KOH → K+ + OH-
Strong bases are completely ionized in solution to form hydroxide ions (OH-). For example, KOH dissolves in water in the reaction
KOH → K+ + OH-
Group 1 and Group 2 metals form strong bases. Examples are
Sodium hydroxide
NaOH
Potassium hydroxide
KOH
Cesium hydroxide
CsOH
Calcium hydroxide
Ca(OH)2

12. Weak bases NH3 + H2O ↔ NH4+ + OH-
A weak base is one which doesn't convert fully into hydroxide ions in solution.
Ammonia (NH3) is a typical weak base. Ammonia itself obviously doesn't contain hydroxide ions, but it reacts with water to produce ammonium ions and hydroxide ions.
NH3 + H2O ↔ NH4+ + OH-
This eqilibrium reaction is reversible, and at any one time about 99% of the ammonia is still present as ammonia molecules. Only about 1% has actually produced hydroxide ions.

13. Don’t confuse…
It is important that you don't confuse the words strong and weak with the terms concentrated and dilute.
It is perfectly possible to have a concentrated solution of a weak acid, or a dilute solution of a strong acid.
For instance, a 12M HF is a concentrated solution of a weak acid, while 0.01M HNO3 is a dilute solution of a strong acid.

14. Electrolytic Properties
Strong acids and strong bases are good conductors of electricity while weak acids and weak bases are not.
Why? Remember that strong acids and strong bases ionize almost 100%, creating many ions which act as mobile charge carriers. Weak acids and weak bases have few ions and few charge carriers.

15. Amphoteric Compounds
The ability of some chemicals to act either as an acid or a base is called amphoterism. Whether an amphoteric chemical acts as an acid or a base depends on what other chemicals happen to be around.
Water is such an amphoteric substance. We say that water has a conjugate base, OH-, and a conjugate acid, H3O+.

16. Amphoteric Compounds
If a base (like ammonia, NH3) is present, water can act as an weak acid and react by donating a proton to that base. In doing so, water is changed into its conjugate base, hydroxide ion.
H2O + NH3 → NH4+ + OH-

17. Amphoteric Compounds
If an acid (like HCl) is present, water can act as a weak base and react by accepting a proton from that acid. In doing so, water is changed into its conjugate acid, hydronium ion.
H2O + HCl → Cl- + H3O+

18. Self-Ionization of Water
The self-ionization of water is another example of water being able to react either as an acid or a base.
The molecules in pure water continuously collide and react with one another. In that reaction, one water molecule can transfer a proton to another water molecule. One water molecule acts as an acid and the other acts as a base.
The solution is neutral because equal quantities of H3O+ and OH- are made.
H2O + H2O → OH- + H3O+

19. More Amphoteric Compounds
Oxides of weakly electropositive metals such as zinc, lead, aluminum, tin, and beryllium (depends on their oxidation state)
Amino acids and proteins, which have amine and carboxylic acid groups
Ammonia (NH3) is another self-ionizable compound
Bicarbonate ion, HCO3-

20. Standard Notation for Acids and Bases
To represent concentrations of ions in moles per liter, the formula of the particular ion or molecule is enclosed in brackets [ ].
For instance [H3O+] means “hydronium ion concentration in moles per liter” or “molar hydronium ion concentration.”
Similarly, [OH-] means “hydroxide ion concentration in moles per liter” or “molar hydroxide ion concentration.”

21. Ion Concentrations as Molarity
Did you notice that we already know a measure of concentration that is expressed in “moles of solute per liter of solution”? Yes, you recognize this as “molarity.”
So both “molar hydronium ion concentration” and “molar hydroxide ion concentration” are expressed as molarity.

22. Standard Values Kw = [H3O+] [OH-]
At room temperature (25 °C), pure water has [H3O+] = 1.0 x 10-7 M and [OH-] = 1.0 x 10-7 M.
The mathematical product of [H3O+] and [OH-] remains constant in water and dilute aqueous solutions at constant temperature.
We call this product the ionization constant of water, Kw, which is expressed as:
Kw = [H3O+] [OH-]
Therefore, at 25 °C , the ionization constant of water has this value:
Kw = [H3O+] [OH-] = 1.0 x 10-14

23. Acidity and Basicity
pH (pouvoir hydrogène or ‘power of hydrogen’) is a measure of the acidity or basicity of an aqueous solution. The molar concentration of hydronium ions in the solution [H3O+] is used in the formula for pH:
pH = -log[H3O+]
(Note: This is log10, not ln, the natural log.)
Solutions with a pH less than 7 are said to be acidic and solutions with a pH greater than 7 are basic or alkaline.
Pure water has a pH very close to 7 (neutral).

24. pH Scale
Pure water

25. Relationship between pH and pOH
Similarly, the concentration of OH ions is expressed as pOH
pOH = -log[OH-]
This leads us to a relationship between pH and pOH at 25 °C:
pH + pOH = 14.0

26. Sample Calculations of pH
What is the pH of a solution that contains a [H3O+] of 1.0 x 10-4 M?
pH = -log[H3O+]
= -log[1.0 x 10-4]
= -(-4)
= 4

27. Sample Calculations of pH
What is the pH of a M solution of hydrochloric acid (HCl)? (Here we use molarity equivalently to concentration.)
pH = - log [.001]
= (.001)
= - (-3.00)
= 3.00
LOG
(-)

28. Sample Calculations of pH
What is the pH of a solution if the hydronium concentration, [H3O+], is x 10-5 M?
pH = - log [H3O+] = - log [3.4 x 10-5 ]
= ( )
= - (- 4.47)
= 4.47
2nd EE
LOG
(-)

29. Sample Calculations of pH
What is the pH of 735 L of a solution holding 0.34 moles of nitric acid (HNO3)?
To find the concentration, we need to find the molarity:
M = 𝒎𝒐𝒍𝒆𝒔 𝒐𝒇 𝒔𝒐𝒍𝒖𝒕𝒆 𝒗𝒐𝒍𝒖𝒎𝒆 𝒐𝒇 𝒔𝒐𝒍𝒖𝒕𝒊𝒐𝒏 (𝑳) = moles nitric acid 735 L
M = M HNO3
Substitute this into the equation for pH:
pH = - log( ) = 3.33

30. Sample Calculations of pH
What is the pH of a 1.0 x 10-3 M NaOH solution at room temperature?
Hmm, NaOH is a strong base, so it will disassociate completely into Na+ and OH- ions. Therefore the concentration of OH- must be the same as the solution’s molarity: [OH-] = 1.0 x 10-3
We use concentration of OH- to find pOH:
pOH = -log[OH-] = -log[1.0 x 10-3]
pOH = 3.0
We know that pH + pOH = 14.0, so
pH = 14.0 – pOH = 14.0 – 3.0 = 11.0
We expected a high pH for a strong base!

31. Calculating [H3O+] from pH
If you know the pH of a substance, you can calculate the ion concentration of [H3O+] from this formula:
[ 𝐻 3 𝑂 + ] = 10 −𝑝𝐻
For a substance with a pH of 7.52, and
using your graphing calculator,
7.52 = E -8
10x ( )
= x 10-8 = [ 𝐻 3 𝑂 + ]
2nd
LOG
(-)

32. Classifying Acids
Most acids used in the laboratory can be classified as either binary acids or oxyacids.
Binary acids are acids that consist of two elements, usually hydrogen and one of the halogens.
Oxyacids are acids that contain hydrogen, oxygen, and a THIRD element, usually a NONMETAL.
‘Acid’ usually refers to a solution in water of one of these special compounds.

33. NAMING ACID RULES: Binary Acids
Anion does NOT contain oxygen (O)
use prefix ‘hydro-’ AND suffix ‘–ic’
added to the root word
Examples: HCl hydrochloric acid
HF hydrofluoric acid
HBr hydrobromic acid

34. NAMING ACID RULES: Oxyacids
Anion DOES contain oxygen (O)
change suffix ‘–ate’ to ‘–ic ‘
(“You ate something icky.”)
Examples: H2SO4 sulfuric acid
HNO3 nitric acid
change suffix ‘–ite’ to ‘–ous’
(“A snake bite is poisonous.”)
Examples: H2PO3 phosphorous acid
HNO2 nitrous acid

35. Naming More Complicated Acids
For example, H2SO5, H2SO4, H2SO3, and H2SO2 are all acids. How do we name them?
Our point of reference is the ‘-ic’ acid made from the ‘-ate’ ion. (Here H2SO4 or sulfuric acid)
Then the acid with one more oxygen than the ‘-ic’ acid is called the per-_________-ic acid. (Persulfuric acid)

36. Naming More Complicated Acids
The acid with one less oxygen than the ‘-ic’ acid is called the ______-ous acid. (This is the ‘-ite’ to ‘-ous’ change.)
If the acid has one less oxygen than the ‘-ous’ acid, it is called the hypo-_______-ous acid.
Examples:
H2SO5 = persulfuric acid HNO4 = pernitric acid
H2SO4 = sulfuric acid HNO3 = nitric acid
H2SO3 = sulfurous acid HNO2 = nitrous acid
H2SO2 = hyposulfurous acid HNO = hyponitrous acid