Transcript Acids & Bases - Independent School District 196

Acids & Bases
Properties of Acids
• Sour taste
• Change color of acid-base indicators (red in pH
paper)
• Some react with active metals to produce
hydrogen gas
Ba(s) + H2SO4(aq)
BaSO4(s) + H2(g)
• Some react with bases to neutralize and form
salt and water
H2SO4 (aq) + 2NaOH(aq)
• Some are electrolytes
Na2SO4 (aq) + 2H2O(l)
Examples of Acids
• Lemons and oranges - citric acid
• Vinegar - 5% by mass acetic acid
• Pop and fertilizer - phosphoric acid
Properties of Bases
• Bitter taste
• Change color of acid-base indicators
(blue in pH paper)
• Dilute aqueous solutions feel slippery
Ex. Soap
• Some react with acids to neutralize and
form salt and water
• Some are electrolytes
Examples of Bases
• Soap - NaOH
• Household cleaners - NH3
• Antacids - Ca(OH)2, Mg(OH)2
Arrhenius Acids
• Acids that increase the concentration of
hydronium (H3O+) in aqueous solutions
HNO3(aq) + H2O(l)
H3O+(aq) + NO3-(aq)
acid
H+ + NO3- + H2O
Why do acids produce H3O+?
• H+ is extremely attracted to the unshared pair of
electrons on the water molecule so it donates itself to
this molecule where it becomes covalently bonded.
The ion formed is known as the hydronium ion (H3O+)
H+
Arrenius Bases
• Bases that increase the concentration of
hydroxide ions (OH-) in aqueous
solutions
H2O
NaOH(s)
Na+(aq) + OH-(aq)
Strength of Acids & Bases
• Strong acids & bases completely ionize in
aqueous solutions
H2SO4 + H2O
NaOH
H3O+ + HSO4Na+ + OH-
• Strong acids & bases are strong electrolytes
• A list of strong acids & bases can be found on
pg. 460-461
• Weak acids & bases only partially break down
into ions when in aqueous solutions
HCN + H2O
NH3 + H2O
H3O+ + CNNH4+ + OH-
• Weak acids & bases are weak electrolytes
• A list of weak acids & bases can be found on
pg. 460-461
Why can we drink H2O?
• Water self ionizes to form equal
concentrations of H3O+ and OHH2O(l) + H2O(l)
H3O+(aq) + OH-(aq)
• A substance is considered “neutral”
when [H3O+] = [OH-]
• [H3O+] concentration = 1.0 x 10-7M
• [OH-] concentration = 1.0 x 10-7 M
When [H3O+] = [OH-]
• If [H3O+] > 1.0 x 10-7 M, the solution is acidic
• If [OH-] > 1.0 x 10-7 M, the solution is basic
• To find the concentration of [H3O+] or [OH-] in
acidic or basic solutions, the following
equation can be used:
1.0 x 10-14 M2 = [H3O+] [OH-]
1.0 x 10-14 M2 = ionization constant for H2O (Kw)
Sample Problem
• A 1.0 x 10-4 M solution on HNO3 has
been prepared for laboratory use.
a. Calculate the [H3O+] of this solution
b. Calculate the [OH-] of this solution
c. Is this solution acidic or basic?
Why?
d. Substitute H2SO4 as the acid. How
would the calculations change?
Sample Problem
• An aqueous 3.8 x 10-3 M NaOH solution has
been prepared for laboratory use.
a. Calculate the [H3O+] of this solution
b. Calculate the [OH-] of this solution
c. Is this solution acidic or basic?
Why?
d. Substitute Ca(OH)2 as the base. How
would the calculations change?
Practice Problems
• Complete practice problems on pg. 484
#1-4
The pH scale
• The pH scale measures the power of the
hydronium ion [H3O+] in a solution
• The scale typically goes from 1-14 (although
it can extend below or above it under extreme
conditions)
• The following equations can be used to
determine the pH or [H3O+] of a solution:
pH = -log [H3O+] [H3O+] = antilog (-pH)
[H3O+] = 1 x 10-pH
pH > 7 basic
pH = 7 neutral
pH < 7 acidic
The pOH scale
• The pOH scale measures the power of the
hydroxide ion [OH-] in a solution
• The scale typically goes from 1-14 (although
it can extend below or above it under extreme
conditions)
• The following equations can be used to
determine the pOH or [OH-] of a solution:
pOH = -log [OH-] [OH-] = antilog (-pOH)
[OH-] = 1 x 10-pOH
pH + pOH = 14
Sample Problems
•
Calculate the pH of each of the
following. Classify as acidic or basic.
a. 1.3 x 10-5 M NaOH
b. 1.0 x 10-4 M HCl
Sample Problems
•
What is the [H3O+] for each of the
following? Classify as acidic or basic.
a. pH = 5.8
b. pOH = 8.9
Sample Problems
•
What is the [OH-] for each of the
following? Classify as acidic or basic.
a. [H3O+] = 9.5 x 10-10 M
b. pOH = 1.3
Practice Problems
• Complete practice problems on
pg. 487 #1
pg. 488 #1-4
pg. 490 #1-4
Strong Acid-Base
Neutralization
• When equal parts of acid and base are
present, neutralization occurs where a
salt and water are formed
HCl(aq) + NaOH(aq)
NaCl(aq) + H2O(l)
Sample Problems
•
•
•
•
H2CO3 + Sr(OH)2
HClO4 + NaOH
HBr + Ba(OH)2
NaHCO3 + H2SO4
Titrations
• When you have a solution with an unknown
concentration, you can find it by reacting it
completely with a solution of known
concentration
• This process is known as titrating
• To perform a titration, an instrument called a
buret can be used to precisely measure
amounts of solution, drop by drop
Titration Termonology
• Equivalence point - the point at which the
known and unknown concentration solutions
are present in chemically equivalent amounts
moles of acid = moles of base
Indicator - a weak acid or base that is added
to the solution with the unknown
concentration before a titration so that it will
change color or “indicate” when in a certain
pH range (table 16-6 on pg. 495 in your text
will show various indicators and their color
ranges)
• End point - the point during a titration
where an indicator changes color
• The 2 most common indicators we will
use in our chemistry class will be:
• Phenolphthalein - turns very pale pink at
a pH of 8-10
• Bromothymol blue - turns pale green at
a pH of 6.2-7.6
Phenolpthalein is clear at pH<8,
pale pink at pH 8-10 and
magenta at pH >10
Bromothymol blue
Practice Titration for
an unknown acid
• 1. Titrate 5.0 of mL of unknown HCl into a 250 mL
erlenmeyer flask - *remember to document the
starting amount and ending amount of acid on the
buret to prevent error
• 2. Add 2 drops of indicator (phenolphthalein) to the
flask - the color of the solution should be clear
• 3. Titrate with .5M NaOH, continuously swirling the
flask, until the solution turns very pale pink for 30
seconds - *remember to document the starting
amount and ending amount of base on the buret
• 4. Mathematically determine the concentration of the
unknown HCl solution by using the following
equation:
Titration Equation
MAVA = MBVB
MA = molarity (mol/L) of acid
VA = volume in L of acid
MB = molarity (mol/L) of base
VB = volume in L of base
molesA = molesB
5. After calculating the molarity of the unknown
acid experimentally, get the theoretical
molarity and calculate % error
Practice titration for an
unknown base
• 1. Titrate 5.0 of mL of unknown NaOH into a 250 mL
erlenmeyer flask - *remember to document the starting
amount and ending amount of base on the buret to prevent
error
• 2. Add 2 drops of indicator (phenolphthalein) to the flask the color of the solution should be magenta
• 3. Titrate with .5M HCl, continuously swirling the flask, until
the solution turns very pale pink for 30 seconds - *remember
to document the starting amount and ending amount of acid
on the buret
• 4. Mathematically determine the concentration of the
unknown NaOH solution by using MAVA = MBVB
• 5. After calculating the molarity of the unknown base
experimentally, get the theoretical molarity and calculate %
error