Antoine Lavoisier, Humphry Davy, and Justus Liebig also made observations regarding acids and bases, but didn't formalize definitions.
Svante Arrhenius
- acids produce H+ ions in aqueous solutions
- bases produce OH- ions in aqueous solutions
- water required, so only allows for aqueous solutions
- only protic acids are allowed; required to produce hydrogen ions
- only hydroxide bases are allowed
Johannes Nicolaus Brønsted - Thomas Martin Lowry
- acids are proton donors
- bases are proton acceptors
- aqueous solutions are permissible
- bases besides hydroxides are permissible
- only protic acids are allowed
Gilbert Newton Lewis
- acids are electron pair acceptors
- bases are electron pair donors
- least restrictive of acid-base definitions
Properties of Acids
- taste sour (don't taste them!)... the word 'acid' comes from the Latin acere, which means 'sour'
- acids change litmus (a blue vegetable dye) from blue to red
- their aqueous (water) solutions conduct electric current (are electrolytes)
- react with bases to form salts and water
- evolve hydrogen gas (H2) upon reaction with an active metal (such as alkali metals, alkaline earth metals, zinc, aluminum)
Properties of Bases
- taste bitter (don't taste them!)
- feel slippery or soapy (don't arbitrarily touch them!)
- bases don't change the color of litmus; they can turn red (acidified) litmus back to blue
- their aqueous (water) solutions conduct and electric current (are electrolytes)
- react with acids to form salts and water
Examples of Common Acids
- 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)
Examples of Common Bases
- detergents
- soap
- lye (NaOH)
- household ammonia (aqueous)
Projects with Acids & Bases
- Apple Browning Lab
- Make Cabbage pH Indicator
- Water into Wine Demonstration
pH Calculations-Acids, Bases, and pH
There are several ways to define acids and bases, but pH only refers to hydrogen ion concentration and is only meaningful when applied to aqueous (water-based) solutions. When water dissociates it yields a hydrogen ion and a hydroxide.
H2O <--> H+ + OH-
When calculating pH, remember that [] refers to molarity, M.
Kw = [H+][OH-] = 1x10-14 at 25°C
for pure water [H+] = [OH-] = 1x10-7
Acidic Solution: [H+] > 1x10-7
Basic Solution: [H+] < 1x10-7
Calculate pH and [H+]
pH = log10[H+]
[H+] = 10-pH
Example:
Calculate the pH for a specific [H+]. Calculate pH given [H+] = 1.4 x 10-5 M
pH = log10[H+]
pH = log10(1.4 x 10-5)
pH = 4.85
Example:
Calculate [H+] from a known pH. Find [H+] if pH = 8.5
[H+] = 10-pH
[H+] = 10-8.5
[H+] = 3.2 x 10-9 M
Binary Acids
A binary compound consists of two elements. Binary acids have the prefix hydro in front of the full name of the nonmetallic element. They have the ending -ic. Examples include hydrochloric and hydrofluoric acid.
- Hydrofluoric Acid - HF
- Hydrochloric Acid - HCl
- Hydrobromic Acid - HBr
- Hydroiodic Acid - HI
- Hydrosulfuric Acid - H2S
Ternary Acids
Ternary acids commonly contain hydrogen, a nonmetal, and oxygen. The name of the most common form of the acid consists of the nonmetal root name with the -ic ending, The acid containing one less oxygen atom than the most common form is designated by the -ous ending. An acid containing one less oxygen atom than the -ous acid has the prefix hypo- and the -ous ending. The acid containing one more oxygen than the most common acid has the per- prefix and the -ic ending.
- Nitric Acid - HNO3
- Nitrous Acid - HNO2
- Hypochlorous Acid - HClO
- Chlorous Acid - HClO2
- Chloric Acid - HClO3
- Perchloric Acid - HClO4
- Sulfuric Acid - H2SO4
- Sulfurous Acid - H2SO3
- Phosphoric Acid - H3PO4
- Phosphorous Acid - H3PO3
- Carbonic Acid - H2CO3
- Acetic Acid - HC2H3O2
- Oxalic Acid - H2C2O4
- Boric Acid - H3BO3
- Silicic Acid - H2SiO3
Bases
- Sodium Hydroxide - NaOH
- Potassium Hydroxide - KOH
- Ammonium Hydroxide - NH4OH
- Calcium Hydroxide - Ca(OH)2
- Magnesium Hydroxide - Mg(OH)2
- Barium Hydroxide - Ba(OH)2
- Aluminum Hydroxide - Al(OH)3
- Ferrous Hydroxide or Iron (II) Hydroxide - Fe(OH)2
- Ferric Hydroxide or Iron (III) Hydroxide - Fe(OH)3
- Zinc Hydroxide - Zn(OH)2
- Lithium Hydroxide - LiOH
What Makes a Strong Acid or Strong Base?
Strong electrolytes are completely dissociated into ions in water. The acid or base molecule does not exist in aqueous solution, only ions. Weak electrolytes are incompletely dissociated.
Strong Acids
Strong acids completely dissociate in water, forming H+ and an anion. There are six strong acids. The others are considered to be weak acids. You should commit the strong acids to memory:
HCl - hydrochloric acid
HNO3 - nitric acid
H2SO4 - sulfuric acid
HBr - hydrobromic acid
HI - hydroiodic acid
HClO4 - perchloric acid
100% dissociation isn't true as solutions become more concentrated. If the acid is 100% dissociated in solutions of 1.0 M or less, it is called strong. Sulfuric acid is considered strong only in its first dissociation step.
H2SO4 -> H+ + HSO4-
Weak Acids
A weak acid only partially dissociates in water to give H+ and the anion. Examples of weak acids include hydrofluoric acid, HF, and acetic acid, CH3COOH. Weak acids include:
Molecules that contain an ionizable proton. A molecule wih a formula starting with H usually is an acid.
Organic acids containing one or more carboxyl group, -COOH. The H is ionizable.
Anions with an ionizable proton. (e.g., HSO4- --> H+ + SO42-)
Cations
transition metal cations
heavy metal cations with high charge
NH4+ dissociates into NH3 + H+
Strong Bases
Strong bases dissociate 100% into the cation and OH- (hydroxide ion). The hydroxides of the Group I and Group II metals usually are considered to be strong bases.
LiOH - lithium hydroxide
NaOH - sodium hydroxide
KOH - potassium hydroxide
RbOH - rubidium hydroxide
CsOH - cesium hydroxide
*Ca(OH)2 - calcium hydroxide
*Sr(OH)2 - strontium hydroxide
*Ba(OH)2 - barium hydroxide
* These bases completely dissociate in solutions of 0.01 M or less. The other bases make solutions of 1.0 M and are 100% dissociated at that concentration. There are other strong bases than those listed, but they are not often encountered.
Weak Bases
Examples of weak bases include ammonia, NH3, and diethylamine, (CH3CH2)2NH.
Most weak bases are anions of weak acids.
Weak bases do not furnish OH- ions by dissociation. Instead, they react with water to generate OH- ions.