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Lactic acid
Found in sour milk.
Citric acid
Found in lemons.
Acetic acid
Found in vinegar.
Phosphoric acid
Found in carbonated beverages.
Malic acid
Found in apples.
Tartaric acid
Found in grape juice.
Ammonia
Used for all types of cleaning.
Sodium hydroxide
Also known as lye
Magnesium hydroxide
Also known as milk of magnesia
Properties of acids
Single replacement reaction
Example: Ba(s) + H2SO4(aq) → Ba SO4 (aq) + H₂ (g).
Binary acid
An acid that contains only 2 different elements: Hydrogen and a more electronegative element.
Oxyacid
An acid that is a compound of hydrogen
Sulfuric acid
Most commonly produced industrial chemical
Nitric Acid
Very volatile
Hydrochloric Acid
Produced by the stomach to aid in digestion
Properties of Bases
Svante Arrhenius
Swedish chemist who theorized that acids and bases produce ions in solution.
Arrhenius acid
A chemical compound that increases the amount of hydrogen ions
Arrhenius base
a substance that increases the concentration of hydroxide ions
Aqueous acids
Water solutions of acids.
Strong acid
ionizes completely in aqueous solution and is a strong electrolyte.
Acid strength
depends on the polarity of the bond between hydrogen and the element it is bonded to and the ease with which that bond can be broken.
Weak acid
weak electrolyte
the amount of hydrogens on a molecule does not affect strength.
Example of weak acid
H3PO4
none of these ionize completely so it is a weak acid.
Organic acids
contain acidic carboxyl group (-COOH) and are weak.
Alkaline
when a base completely dissociates in water to yield aqueous (OH) ions.
Example of strong base
NaOH.
Sodium
an alkali metal
the group gets its name because when the elements combine with hydroxides they form alkaline solutions.
Strong bases
strong electrolytes.
Alkalinity
depends on the concentration of hydroxide ions in solution
Example of weak base
Ammonia (NH3) is highly soluble but a weak electrolyte.
Bronsted-Lowery acid
a molecule or ion that is a proton donor (H+).
Example of Bronsted-Lowery acid
Hydrochloric acid
Bronsted-Lowery base
a molecule or ion that is a proton acceptor.
Bronsted-Lowery reaction
proton transferred from one reactant (the acid) to another reactant (the base).
Monoprotic acid
an acid that can donate only one proton (hydrogen ion) per molecule.
Examples of monoprotic acids
HCl
Polyprotic acid
an acid that can donate more than one proton per molecule.
Example of diprotic acid
H2SO4 (sulfuric acid).
Example of triprotic acid
H3PO4 (phosphoric acid).
Conjugate base
the species that remains after a Bronsted-Lowery acid has given up a proton.
Conjugate acid
the species that is formed when a Bronsted-Lowery base gains a proton.
Strength of conjugate acids and bases
depends on the strength of the acids and bases involved. The stronger an acid is
the stronger a base is
the weaker its conjugate acid.
Amphoteric compounds
any species that can react as either an acid or a base.
Neutralization
the reaction of hydronium ions and hydroxide ions to form water molecules.
Salt
an ionic compound composed of a cation from a base and an anion from an acid.
Spectator ions
ions that do not participate in the overall reaction.
Net ionic equation
an equation that shows only the species that actually change during the reaction.
Acid rain
precipitation that occurs when gases such as NO
Calcium carbonate reaction with acid
CaCO3 (s) + 2H3O+ (aq) → Ca²+ (aq) + CO2 (g) + 3H₂O (l)
Hydriodic acid
HI
Iodide ion
I-
Perchloric acid
HClO4
Perchlorate ion
ClO4-
Hydrobromic acid
HBr
Bromide ion
Br-
Chloride ion
Cl-
Hydrogen sulfate ion
HSO4-
Nitrate ion
NO3-
Acetate ion
CH3COO-
Carbonic acid
H2CO3
Hydrogen carbonate ion
HCO3-
Hydrosulfuric acid
H2S
Hydrosulfide ion
HS-
Dihydrogen phosphate ion
H2PO4-
Hydrogen phosphate ion
HPO4^2-
Hypochlorous acid
HClO
Hypochlorite ion
ClO-
Ammonium ion
NH4+
Hydroxide ion
OH-
Amide ion
NH2-
Hydride ion
H-