Name: _______________________________

Unit 11: Reactivity: Acid-Base, Precipitation, Oxidation-Reduction(Redox) ( 3 categories of reactions )
Properties of Acids: Tart or _Sour_; React with certain metals_to produce hydrogen gas. React with bases to
produce water and a salt; Turns blue litmus to the color __RED_____

Electrolytes = __conduct_______ electricity; Acids – vinegar, carbonated drinks, citrus fruit

Properties of Bases: Bitter taste; Feel _slippery; React with acids to form water and a salt

Turns red litmus to the color blue; Electrolytes = __conduct___ electricity; Bases – antacid tablets, household
cleaning agents

Ions in Solution: ___Acidic_______– contain more H+ than OH-

_Basic_________– contain more OH- than H+ ; _Neutral____– contain equal amounts of H+ and OH-

Naming Acids - reminder - I ate something icky. I bite something delicious.

Naming Bases - reminder - A base is a compound that produces hydroxide ions, (OH-), when dissolved in water.

Bases are named the same way any other ionic compound is named. NaOH – sodium hydroxide.

Arrhenius Acids and Bases:

Svante Arrhenius (Swedish chemist) (1859-1927); 1890’s – formulated first useful theory of acids and bases.

Acid: contain H+ ; the dissolved compound pr0duces a hydrated hydrogen ion (H3O+(aq) , hydronium) and an
anion.

Base(alkali): contains at least one OH- ion that dissociates when dissolved in water. Dissolved compound
produces hydroxide ion and a cation.

Reaction produces H2O and a dissolved salt. This process is known as Neutralization.

Bronsted-Lowry Acids and Bases: 1920’s – definitions of acids and bases expanded independently by –
Danish chemist Johannes Bronsted and English chemist Thomas Martin Lowry .

More broad definition of acids/bases

Acids- hydrogen ion, or proton, donors; Bases- hydrogen ion, or proton, acceptors

A conjugate acid and conjugate base are formed.

Base gains hydrogen ion; particle formed is conjugate acid.

The particle that remains when an acid has donated a hydrogen ion is the conjugate base.

Example: when HCl gas dissolves and ionizes in water, HCl is the proton donor (acid), water is the proton
acceptor (base), the hydronium ion is the conjugate acid, and the chloride ion is the conjugate base.
HCl (g) + H2O (l)
acid
base
→
H3O+ (aq)
conjugate acid
+ Cl- (aq)
conjugate base
Name: _______________________________

Neutralization Reactions: acid + base  a salt + water

Neutralization reactions are a type of double replacement reactions.

Water: Water can act as an acid or a base;

HCl + NaOH  NaCl + H2O
H2O as an acid donates one H+ to become OH-
H2O as a base accepts one H+ to become H3O+
This is called the Hydronium ion. (H3O+)

Neutral Solution—Any aqueous solution in which the concentration of H+, [H+], and the concentration of OH-,
[OH-], are equal.

[H+] = hydrogen ion concentration

[OH-] = hydroxide ion concentration
In pure water at 25 oC, [H+] = [OH-] = 1.0 x 10-7 M.
In any aqueous solution, [H+] and [OH-] are interdependent. That is, when [H+] increases, [OH-] must decrease, and vice
versa.

Ion-product Constant for Water (Kw)

The product of the concentrations of the hydrogen ions and hydroxide ions in water is called the ion-product
constant for water (Kw).

Kw =

Kw = (1.0 x 10-7M) x (1.0 x 10-7M) = 1.0 x 10-14 M

If you know one, you can calculate the other. pH scale is derived from the value of Kw. Brackets, [ ], denote
oncentration of”

The pH concept:
pH is the French abbreviation for “pouvier d’hydrogen.” In English, this translates to the “power of hydrogen.”

The pH of a solution is the negative logarithm of the hydrogen-ion concentration.

Pure water, pH = 7, is said to be neutral.

The formula is:
[H+]
x
[OH-]
= 1.0 x 10-14 M
pH = -log[H+]
Since Kw = 1 x 10-14, for any solution:
pH = -log (1 x 10 -7) = 7
Likewise…………. pOH = -log[OH-]
pH + pOH = 14 so, pOH = 14 – pH and pH = 14 - pOH

Each unit of pH represents a power of 10.

Example: Baking soda has pH of 8.4 and Milk has pH of 6.4. This difference is 2 pH units.
Since each unit represents a power of 10, the [H+] in milk is 102 (or 100) times the [H+] in baking soda.

Strength and Dissociation :

Strong & weak acids are determined by how much they dissolve into ions in water.

Strong acids – completely ionize, higher conductivity, lower pH than same concentration of weak acid.
examples: HCl and H2SO4
Name: _______________________________
Weak acids - only partially ionize, lower conductivity.
examples: carbonic acid, acetic acid
Strong base - higher conductivity, higher pH than same concentration of a weak base, example: NaOH
Weak base - lower conductivity. examples: ammonia & sodium bicarbonate
Acid-Base Indicators:

An indicator is an acid or a base that undergoes dissociation in a known pH range.

An indicator is useful for measure of pH because its acid form and base form have different colors in solution.

Ex: litmus paper, pH paper, phenolphthalein
Acid-Base Indicator Table (See pg. 590)
pH Meters:

A __pH Meter__ is used to make rapid, accurate pH measurements.

If a pH meter is connected to a computer or chart recorder, it can be used to make a continuous recording of pH
changes.

Most pH meters are accurate to within 0.01 pH unit of the true pH.
The _strength__ of an acid or a base tells you the degree of ionization.

Strong acids & bases break down into __all_ ions. (Completely ionized, or dissociated)


Ex: Strong acid – hydrochloric acid, sulfuric acid. Strong base – Sodium hydroxide.
Weak acids & bases break down into just a few ions (Partially ionized, or dissociated)
Ex: Weak acid – acetic acid, carbonic acid. Weak base – Sodium bicarbonate (baking soda)
Titration: Lab technique used to experimentally determine the molarity of a solution.

A titration is a carefully monitored addition of one solution into another solution.

Often done by the use of a burette, which allows for accurate measurement needed to neutralize the other
solution. An indicator or pH meter is needed to identify at what point the neutralization process is complete.

(ex: Phenolphthalein - colorless to pink around pH=9)

Oxidation-Reduction ………………Also known as redox reactions

Very common in nature – rusting and other types of corrosion are typical examples

A chemical reaction in which one reactant loses electrons and another gains electrons.

Example:
Fe →
Fe+2 + 2e- (Iron loses e-s.)
O2 + 4e- → 2O-2 (Oxygen gains e-s.)
Put together:
2Fe + O2
→
2FeO
Note: In order to be balanced, have total of 4 electrons lost by iron and 4 gained by oxygen.
Name: _______________________________

Redox (cont.)

The reactant being oxidized (losing electrons) is called the reducing agent.

The reactant that is reduced (gaining electrons) is called the oxidizing agent.

Typically, redox reactions happen in water or moist surroundings.

Note: ALL batteries produce electricity via redox reactions. Regular batteries eventually consume the reactants.
Rechargable batteries reverse the redox reaction by adding electricity (electrons), so they last longer.
Remember to use Activity Series of Metals for info, too!
Precipitation:

Precipitation reaction may occur when two solutions of IONIC substances are mixed.

If one or both pairs of ions in the combined solutions form ionic compounds with very little to no solubility in
water (unable to dissolve or very little), then a new, insoluble ionic solid will precipitate out.

Note: this is NOT the same as the physical change of crystallization seen in some solutions due to temperature
changes.