Transcript Acids, Bases, and Neutralization
Acids, Bases, and Neutralization
Naming Acids
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Acids - All have H+ , so the name is based on the anion (- part) Anion Rule Example ending HCl has Chloride =Hydrochloric acid -ide Hydro(stem)ic acid HBr has bromide = Hydrobromic acid -ite -ate (stem)ous acid (stem)ic acid HNO 2 has nitrite = nitrous acid HClO 2 has chlorite = chlorous acid H 3 PO 4 acid has phosphate = phosphoric H 2 SO 4 has sulfate = sulfuric acid HNO 3 has nitrate = nitric acid *S compounds use (sulf) as the root * P compounds use (phosphor) as the root
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Formulas of Acids
Use H+ as the cation. Determine the anion from the name. Combine enough of each to balance out the charge.
Example
hydronitric acid nitride Phosphorous acid carbonic acid anion phosphite carbonate ions H + N -3 H + PO 3 -3 H + CO 3 -2 formula H 3 N H 3 PO 3 H 2 CO 3
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Bases are named like any other ionic compound: Cation Anion
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Formulas are “ “ “ “ “
Bases
• • • Naming Bases = like any ionic compound • Cation anion (change ending to –ide if not a polyatomic ion) • • NaOH = ____________________ Ca(OH) 2 =______________________ Base formulas – like any ionic compound EX: Magnesium hydroxide Mg +2 OH = Mg (OH) 2
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Neutralization Reactions
Reaction in which an acid and a base react in an aqueous solution to produce a salt and H 2 O (Double Replacement reaction)
Steps
1.
Write acid and base formulas 2.
3.
Write HOH(water) as product Combine remaining cation and anion using subunits to make neutral compound (salt) 4.
Balance
Example:
HNO 3 H 2 SO 4 + KOH 2HCl + Mg(OH) 2 + 2NH 4 OH
Properties of Acids
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Turns Litmus red
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Have pH less than 7
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Taste sour
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Reacts with bases in a neutralization reaction
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Reacts with metals to produce H 2 Ex: Lemon juice
Properties of Bases
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Turns Litmus blue
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Have a pH greater than 7
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Taste bitter
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Feel slippery
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Neutralize acids
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Ex: Soap
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Acids and bases are both strong or weak electrolytes (conduct electricity)
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Electrolytes = dissociate (break apart into ions) when dissolved
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Strong = completely
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Weak = partially
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Non = not at all
Types of Acids/Bases
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Arrhenius Acids/Bases 1884
• Acids release H + into water • Bases release OH into water Ex: Acid: HCl (aq) Base: NaOH (aq) H + Na + (aq) + Cl (aq) (aq) + OH (aq) • • • Monoprotic = acid that contains 1 proton (H + ) • Ex.) HBr, HC 2 H 3 O 2 Diprotic = 2 protons, 2H + Ex.) H 2 SO 4 Triprotic = 3 protons, 3H + Ex.) H 3 PO 4
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Bronsted-Lowry Acids/Bases 1923 - 2 independent scientists
Acid = H + Donor Base = H + Acceptor • Ex: NH 3 + HOH ↔ NH 4 + + OH • Amphoteric = substance that acts as both an acid and a base NH 3 + HOH ↔ NH 4 + + OH • HCl + HOH ↔ H 3 O + + Cl • •
Lewis Acids/Bases
Acid = accepts a pair of e Base = donates a pair of e-
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Strong Acids (Memorize) HBr = Hydrobromic acid HCl = Hydrochloric acid HI = Hydriodic acid HNO 3 HClO 4 = Nitric acid = perchloric acid H 2 SO 4 = Sulfuric acid All other acids are weak
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Strong acids dissociate completely HCl
H + + Cl (break apart 100%)
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Weak Acids Dissociate (ionize) only slightly HC 2 H 3 O 2 (99%) + H 2 O ↔ H 3 O + C 2 H 3 O 3 (1%)
Strong Bases (memorize) KOH Ba(OH) 2 Ca(OH) 2 NaOH LiOH •
Conjugate Acids/Bases
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****Products*** (right side of arrow only)
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Conjugate acid = particle formed when base gains a H +
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Conjugate base = particle formed when acid has lost/donated H + EX: NH 3 + H 2 O
NH 4 + + OH *** Base becomes conj. Acid *** Acid becomes conj. Base EX: HCl + H 2 O
H 3 O + + Cl -
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19.2: H
+
& Acidity
H + from water (Self ionization) : H 2 O(
l
)
H + (
aq
) + OH (
aq
)
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H + joined to H 2 O as hydronium (H 3 O + ) [OH ] = 1 X 10 -7 M [ ] = molarity concentration
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[H + ] = 1 X 10 [OH ] = [H + -7 M ] in pure water, so it is a neutral solution
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K w K w = water dissociation constant = [H + ] [OH ] = 1 X 10 Acidic Solution [H + -14 ] > [OH ]
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Basic Solution [H + ] < [OH ] Alkaline Solution Neutral solution [H + ] = [OH ]
pH
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Soren Sorensen
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Easier way to express [H + ], from 0-14.
pH = - log [H
+
]
Acidic Solution Neutral Solution Basic Solution pH <7.0
pH = 7.0
pH >7.0
[H + ] > 1 X 10 -7 [H + ] =1 X 10 -7 [H + ] <1 X 10 -7
pOH
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pOH expresses [OH ] instead of [H + ]
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pOH = - log [OH
-
]
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pH + pOH = 14
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Examples
Find pH if [H + ] = 1 X 10 -9
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Find [H + ] if pH =6
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Find pOH if pH = 6.8
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Find pOH if [H +
] = 1 X 10
-9
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Find [OH ] if [H +
] = 1 X 10
-9
Indicators = pH demo
Titrations:
Method of determining the concentration of an acid or a base.
Steps for calculations 1.
Write a balanced neutralization equation.
2.
Calculate the moles of known solution using : mol = M*L 3.
4.
Use the balanced equation to find the moles of unknown starting w/ moles of known. Set up a proportion, dividing by the coefficients.
Calculate the concentration of unknown, using calculated moles and measured volume. M=mol/L EX. What is the molarity of phosphoric acid, if 15.0 mL of the solution is titrated by 38.5 mL of 0.15 M sodium hydroxide?
Another way
EX: How many mL of 0.45M H 2 SO 4 KOH to make a neutral solution?
must be added to 25.0 mL of 1.00 M
Titration Graph
K
a
, K
b
, K
eq
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Strong Acids & Bases = dissociate 100%
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(all of it ionizes or breaks apart into ions) H 2 SO 4
2 H + + SO 4 -2 0% 100% Weak acids & Bases = only dissociate partially
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H 3 N
3H + + N -2 95% 5%
Equilibrium
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Equilibrium = rate of forward and reverse reactions are equal (amounts of products and reactants do not change)
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When we have equilibrium, we can write, K a , K b , K eq K eq = Equilbrium constant K eq = [Products] coeff [Reactants] coeff
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K
a
& K
b K a = acid dissociation constant. A measure of how much a weak acid breaks into ions. K eq
for a weak acid:
H 2 X
2H + + X -2 K a =
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Ex: H 3 PO 4
3H + + PO 4 -3 K a =
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Larger K a value = stronger weak acid
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K b = base dissociation constant. A measure of how much a weak base breaks into ions. K eq for a weak base Ex: Be(OH) 2
Be +2 + 2OH K b =
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Larger K b value = stronger weak base