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Acids/Bases/Salts
Properties
Common Acids
Lactic sour milk
Acetic vinegar
Phosphoric tart taste in soda
Citric citrus fruits
Malic apples
Tartaric grapes
Formic ant bites
Common Base
Ammonia window cleaner
Sodium hydroxide lye (drain and oven cleaners)
Sodium Bicarbonate Baking soda
antacid
Milk of magnesia antacid
Properties
electrolytes
turn litmus red
sour taste
react with metals to form H2 gas
slippery feel
turn litmus blue
bitter taste
vinegar, soda, apples, citrus fruits
ammonia, lye, antacid, baking soda
electrolytes
pH less than 7 pH more than 7
Acid nomenclature –naming acids
Binary acids – is an acid that contains only two different elements.
Acids are composed of hydrogen (H+) followed by an anion (negative ion).
Oxyacids- is an acid that is a compound of hydrogen, oxygen and third element (usually a non-metal)
If the acid formula contains oxygen in the anion, such as in H2SO4, it is known as an oxyacid.
3 Rules To Naming Acids
If H + anion ending in –ide: Acid name is “hydro_____ic acid” Take the root from the anion name and
fill in the blank.
Acid Naming Example
Example: HCl • Cl is the anion, its name is chloride• Name of acid is: hydrochloric acid
Example: HF• F is the anion, its name is fluoride• Name of acid is: hydrofluoric acid
3 Rules To Naming Acids
H + anion ending in –ate: Acid name is “_____ic acid” Take the root from the anion name and fill in
the blank. “What I ATE was ICky”
Acid Naming Example
Example: HNO3
NO3 1- is the anion, its name is
nitrate Name of acid is: nitric acid
Example: H2CO3
CO3 2- is the anion, its name is carbonate Name of acid is: carbonic acid
Exceptions
Sulfate (SO4 2-) Root is not sulf, but sulfur
• Sulfuric acid
Phosphate (PO4 3-) Root is not phosph, but phosphor
• Phosphoric acid
3 Rules To Naming Acids
H + anion ending in –ite: Acid name is “_____ous acid” Take the root from the anion name and fill in
the blank. “Don’t bITE; it’s infectiOUS”
Acid Naming Example
Example: HNO2
NO2 1- is the anion, its name is nitrite Name of acid is: nitrous acid
Example: HClO2
ClO2 1- is the anion, its name is chlorite Name of acid is: chlorous acid
Strength of Acids
Strong Acid
- is one that ionizes completely in aqueous solution.
- completely dissociates.
- a strong acid is a strong electrolyte.
- increases with increasing polarity and decreasing bond energy.
Strength of Acids
Weak acids
- are weak electrolytes
- it contains hydronium ions, anions and dissolved acid molecules
- Organic acids (acidic carboxyl group) -COOH
Aqueous solutions for base
Most bases are ionic compounds containing metal cations and the hydroxide anion.
When a base completely dissociates in water to yield aqueous OH- ions, the solution is called alkaline.
NaOH (s) water Na+(aq) + OH-
(aq)
Aqueous solutions for base
Not all bases are ionic compounds.
Ammonia is one example because it produces hydroxide ions when it reacts with water molecules.
NH3 (g) + H20 (l) NH4+ (aq) + OH-
(aq)
Strength of bases
The strength of the base depends on the extent to which it dissociates to its ions.
Strong bases are strong electrolytes
Bases that are not very soluble do not produce a large number of hydroxide ions when added to water
Base Naming Example
NaOH name of base: sodium hydroxide
Mg(OH)2
name of base: magnesium hydroxide Fe(OH)2
name of base: iron (II) hydroxide
pH and pH scale
0
7INCREASING
ACIDITY NEUTRALINCREASING
BASICITY
14
Whether or not a solution is acidic, basic, or neutral depends on the balance of H+ and OH- ions: Neutral: [H+] = [OH-]Acid: [H+] > [OH-]Base: [H+] < [OH-]
pH and pOH Calculations
pH is the negative base 10 logarithm of the hydrogen ion concentration:
pH = - log10 [H+]
pH is the expression of the acidity or alkalinity of a solution in terms of its hydronium ion concentration.
pH = - log [H3O+]The sum of the pH and the pOH always equals 14.
pH + pOH = 14
Example
Calculate the pH, if the
[H3O] = 2.4 X 10-6 M
pH = - log [H3O+]
= - log(2.4 X 10-6)
= -(-5.6)
= 5.6 Find the pH, the pOH = 5.3
pH + pOH = 14
pH = 14 - 5.3
pH = 8.7
pH calculations
Use the reverse of the equation to calculate the [H3O+] when pH is known.
[H3O+] = antilog (-pH)
= 10-pH
Use an identical equation to calculate pOH. pOH = - log [OH-]
Calculate the [H3O+], if the pH is 4.71. [H3O+] = antilog (-pH)
= antilog (-4.71) = 1.95 X 10-5 M
Acid- Base Theories
Type Acid Base
Arrhenius(Traditional)
H+ or H30 +
producer
OH –
producer
Bronsted-Lowry proton donor proton acceptor
Lewis electron –pair acceptor
electron- pairdonor
Arrhenius Acids and BasesArrhenius acid – is a chemical compound
that increases the concentration of hydrogen ions H+, in an aqueous solution.
Arrhenius base – is a substance that increases the concentration of hydroxide ions OH-, in aqueous solution.
Acid-Bases Theories
Acid-Bases TheoriesBronsted-Lowry acid
-Is a molecule or an ion that is a proton donor (H+)
- Donates proton to water
HCl + H2O H3O+ + Cl-
HCl + NH3 NH4+ + Cl –
In both reaction HCl is a Bronsted-Lowry acid
Bronsted-Lowry Acid
Water could be a Bronsted-Lowry acid also.
H2O (l) + NH3 (g) NH4+
(aq)+ OH-(aq)
The water donates a hydrogen ion (proton) to the ammonia molecule.
Bronsted-Lowry Base
A molecule or ion that is a proton acceptor (H+)
H2O (l) + NH3 (g) NH4+
(aq)+ OH-(aq)
Ammonium (NH4) is a Bronsted-Lowry Base because it accepts the proton from the water.
Bronsted-Lowry Acid-Base Reaction
Protons are transferred from one reactant (the acid) to another (base).
H2O (l) + NH3 (g) NH4+
(aq)+ OH-(aq
Bronsted-Lowry Acid-Base Monoprotic Acids – an acid that can donate
only one proton (HCl, HNO3)
Polyprotic Acids – an acid that can donate more than one proton per molecule.
- Diprotic – can donate two protons per molecule (Sulfuric acid, H2SO4)
- Triprotic – can donate three protons per molecule(Phosphoric acid, H3PO4)
Acid –Base Reaction
The Bronsted-Lowry definition s of acids and bases provide the basis for studying proton (H+) transfer reaction.
Suppose a Bronsted-Lowry acid gives up a proton, the remaining ion or molecule can re-accept that proton and can act as a base- a conjugate base.
Acid –Base Reaction
Conjugate base- the ion or molecule that remains after a Bronsted-Lowry acid has given up a proton is the conjugate base of that acid.
Conjugate acid- the ion or molecule that is formed when a Bronsted-Lowry base gains a proton is the conjugate acid of that base.
Acid –Base Reaction
HF + H2O F- + H3O+
Acid BaseConjugate Base
Conjugate Acid
The species remaining after a Brønsted-Lowry acid gives up its proton is the conjugate base of that acid: Take off one H from the acid.
The species remaining after a Brønsted-Lowry base accepts its proton is the conjugate acid of that base: Take off one H. Add an H to the base.
Acid –Base ReactionHCO3
- (aq) + H2O (l) H2CO3 (aq) + OH-
(aq)
base acid conjugate acid
conjugate base
(proton acceptor)
HF (aq) + H2O (l) F- (aq) + H3O+
(aq)
acid baseconjugate base
conjugate acid
(proton donor)
Strength of conjugate acids and bases
The stronger an acid is, the weaker its conjugate base; the stronger a base is, the weaker its conjugate acid.
The weaker an acid is, the stronger its conjugate base; the weaker a base is, the stronger its conjugate acid.
Lewis Acids and Base
A Lewis acid is an atom, ion, or a molecule that accepts an electron pair to form a covalent bond.
A Lewis base is an atom, ion or a molecule that donates an electron pair to form a covalent bond.
Lewis Acids and Base
H+ (aq) + : NH3
(aq) [H-NH3
+ ] (aq)
H+ (aq) + : NH3
(aq) [NH4
+ ] (aq)
A bare proton (H+) is a Lewis acid in reactions in which it forms a covalent bond.
Lewis Acids and Base . .
BF3 (aq) + : F : - (aq) BF4 – (aq)
. .
An anion is a Lewis Base in a reaction in which it forms a covalent bond by donating an electron pair.
Indicators Chemical dyes that change color as pH
changes. Different indicators change colors at
different pH levels choose an indicator that will show a color
change at the pH that you are interested in. Indicators can be on a strip of paper
called pH or litmus paper Other indicators can be added to the
solution directly. Some indicators change color more than once
and can be added to solutions so that we can see what is happening over time.
Acid/Bases/Salts
Neutralization/Titrations
Neutralization
Chemical reaction between an acid and a base.
Products are a salt (ionic compound) and water.
Neutralization
ACID + BASE SALT + WATER
HCl + NaOH NaCl + H2O
HC2H3O2 + NaOH NaC2H3O2 + H2O
Salts can be neutral, acidic, or basic.
Neutralization does not mean pH = 7.
weak
strong strong
strong
neutral
basic
Neutralization
HCl + NaOH
H2SO4 + KOH
HNO3 + Ca(OH)2
NaCl + H2O
K2SO4 + H2O
Ca(NO3)2 + H2O
-1 +1
-2 +1
-1+2
Hydrocholoric acid Sodium
hydroxide
Sodium chloride
Water
Sulfuric acid Potassium hydroxide
Potassium sulfate Water
Nitric acid Calcium hydroxide
Calcium nitrate
water
2 2
2 2
Titration
Titration Analytical method in
which a standard solution is used to determine the concentration of an unknown solution.
standard solution
unknown solution
End Point – point at which an indicator changes
color during a titration Equivalence point
Point at which equal amounts of H3O+ and OH- have been added.
when mole ratio exactly equals mole ratio required by reaction
Determined by…• indicator color change
Titration
• dramatic change in pH
Titration
moles H3O+ = moles OH-
MV n = MV nM: MolarityV: volumen: # of H+ ions in the acid
or OH- ions in the base
Titration
42.5 mL of 1.3M KOH are required to neutralize 50.0 mL of H2SO4. Find the molarity of H2SO4.
H3O+
M = ?V = 50.0 mLn = 2
OH-
M = 1.3MV = 42.5 mLn = 1
MV# = MV#M(50.0mL)(2)
=(1.3M)(42.5mL)(1)
M = 0.55M H2SO4