Transcript ACIDS AND BASES - St. Dominic High School


Acidity is a measure of the hydrogen
(hydronium) ion concentration of a solution
(Acid).

Alkalinity is a measure of the hydroxide ion
concentration of a solution (Base).

Arrhenius Acid is a substance that produces
hydronium ions (H3O+) as the only positive
ions when dissolved in water.

Arrhenius Base is a substance that produces
hydroxide ions (OH-) as the only negative
ions when dissolved in water.

Neutralization is a reaction between an
acid and a base to produce salt and
water.
Acid + Base
HNO3 + NaOH
Salt + Water
NaNO3 + H2O

Electrolyte is a substance whose water
solution conducts an electric current (H3O+
or OH-).BASE

Salt is a product (other than water) of a
neutralization reaction; an ionic substance.
- gives the spectator ions; ex. Na+ and Cl-
According to Bronsted and Lowry
Acid: proton (H+) donor
H+ proton – remove electron from H leaving
one proton
Acid is a substance that gives up H+ ions to
another substance
Base: proton (H+) acceptor
Base is a substance that takes an H+ ion from
another substance
CONJUGATE ACID-BASE PAIRS:
A conjugate pair refers to acids and bases
with common features. These common
features are the equal loss/gain of protons
between the pairs. Conjugate acids and
conjugate bases are characterized as the
acids and bases that lose or gain
protons. In an acid-base reaction, an acid
plus a base reacts to form a conjugate
base plus a conjugate acid

Acid + Base→Conjugate Base + Conjugate Acid

The conjugate acid of a base is formed
when the base gains a proton. Refer to
the following equation:
Examples:
HCl + H2O  H3O+ + Cl-
NH3 +H2O  NH4++ OH-
HC2H3O2 + H2O  H3O+ + C2H3O2-
Acid (H+ or H3O+) & Metal
- to react with an acid, the metal must be
higher on Table J than H2
Ca + 2HCl
H2 + CaCl2
Cu + H3PO4
No Reaction


ACIDS:
1. Sour Taste
2. Aqueous solution electrolytes
3. React with bases to produce water and salt
4. React with certain metals to produce H2 gas
5. Cause acid-base indicators to change
color(turn blue litmus red)
6. pH less than 7
7. Acids can be formed by a reaction of gaseous
oxides with water

1.
2.
3.
4.
5.
6.
7.
BASES
Bitter taste
Aqueous solutions are electrolytes
Slippery feel
React with acids to form salt and water
Cause acid-base indicators to change colors
(turn red litmus blue)
pH greater than 7
Formed when Group 1 and 2 metals react
with H2O, H is released too
HCl – dangerous acid
 Citric Acid – in fruits
 Boric Acid – eye washing

These strong acids and bases are STRONG
ELECTROLYTES – GOOD CONDUCTORS of
electricity
RELATIVE STRENGTHS OF ACIDS AND BASES
 Strength is determined by the position of
the “dissociation” equilibrium

Strong Acids (H+ or H3O+)
-
Found at the top of Table K
-
Low pH number (1, 2 or 3)
-
The more oxygen present in the
polyatomic ion of an oxyacid (acid
containing oxygen), the STRONGER its
acid within that group is

Strong Bases (OH-)
-
Found at the top of Table L
-
High pH number (12, 13 or 14)
-
Hydroxides or oxides of group 1 and 3
metals (except Mg and Be)
-
Those that are very soluble are very
strong
WEAK ACIDS AND BASES

Majority of acids and bases are weak –
they do not ionize much
BINARY ACIDS: composed of H and one
other elements (TABLE K)
Ex) HCl - hydrogen chloride
Name begins with hydro followed by the name
of other element with modified ending – ic
Ex) HCl – hydrochloric acid
TERNARY ACIDS: Have polyatomic ions
containing O – table E
1.Anion suffixes – ate and – ite usually
replaced by –ic and –ous respectively
-ate to –ic
-ite to –ous
Ex) HNO3 -> NITRIC ACID
H2SO4 -> SULFURIC ACID
BASES:
-
EX)
Positive ion name is not modified and
then of the base ends with hydroxide
Ca(OH)2  Calcium hydroxide
It was found that no matter how pure
water is, it still conducts a minute current.
This proves that water self-ionizes
 Since water will dissociate with itself to a
slight extent only about 2 billion water
molecules are ionized at any instant
 The equilibrium expression used here is
referred to as the autoionization constant
for water – Kw

In pure water or dilute solutions, the
concentration of water can be
considered to be a constant value, so
we include that with the equilibrium
expression and write it as:
 Kw = [H+] [OH-] = 1.00 x 10-14


Knowing this value allows us to calculate
the ion concentration for OH- and H+ for
various situations
[OH-] = [H+] solution is neutral (in pure
water, each of these is 1.0 x 107
 [OH-] > [H+] solution is basic
 [OH-] < [H+] solution is acidic

Used to designate the [H+] in most
aqueous solutions where [H+] is small
 pH = -log [H+]
 pOH = -log [OH+]
 pH + pOH = 14

0 = strongly acidic
 7 = neutral
 14 = strongly basic

LOGARITHMIC SCALE
 Each change of a pH unit signifies a
tenfold change in the concentration of the
hydrogen ion
-
 Ex)
Concentration of [H+] is 10x greater in
a solution with a pH = 5 than pH = 6
As concentration of H+ increases the
concentration of OH- decreases
- Acidity: strength of acid in H+ ions
- Alkalinity: strength of base in OH- ions
-
1.
pH probes – electrodes that detect
electrical conductivity
2.
Acid –base indicators and narrowing
down pH using multiple indicators (mixture
of indicators gives great range of colors,
pH paper)
-
Chemicals that have certain colors depending
on pH
pH paper to compare results to chart
TABLE M
Ex) A solution yields the following results when
tested with various indicators
Methyl Orange = yellow
Phenolphthalein = clear
Bromcresol green = blue
Thymol blue = yellow
AIM: HOW CAN WE DETERMINE THE
CONCENTRATION OF AN ACID OR BASE?

Controlled process of acid-base neutralization
used to determine the concentration of an
acid or a base

Endpoint: the pH at which an indicator that
has been added to a titration set up turns
color

Equivalence Point: the point at which the
titrated solution has a pH of 7
[OH-] = [H+] – can be detected using probes
1 mole H+ neutralizes 1 mole of OHMoles of H+ = Moles of OHM = mol/L
mol = M x L = M x V


Macid x Vacid = Mbase x Vbase
#H(MaVa) = #OH(MbVb)
Want the number of moles
#H(molesa) = #OH(molesb)
1.
How many moles of LiOH are needed to
exactly neutralize 2 moles of H2SO4?
2.
How many moles of H2SO4 are needed to
exactly neutralize 5.0 moles of NaOH?
3. How many moles of HCl are needed to
neutralize 0.10L of 2.0M solution of NaOH?
4. How many moles of NaOH are needed to
neutralize 0.10L of 0.20M H2SO4 solution?
5. If it takes 15.0mL of 0.40M NaOH to
neutralize 5.0mL of HCl, what is the molar
concentration of the HCl solution?
6. If it takes 10.0mL of 2.0M H2SO4 to neutralize
30.0mL of KOH, what is the molar
concentration of KOH?
7. How many mL of 2.0M H2SO4 are
required to neutralize 30.0mL of 1.0M
NaOH solution?
8. How many mL of 0.10M Ca(OH)2 are
required to neutralize 25.0mL of 0.50M
HNO3 solution?

A titration curve is drawn by plotting
data attained during a titration, titrant
volume on the x-axis and pH on the yaxis. The titration curve serves to profile
the unknown solution. In the shape of the
curve lies much chemistry and an
interesting summary of what we have
learned so far about acids and bases.

The endpoint is when the indicator color
changes – if you picked the right
indicator the equivalence point and the
endpoint will occur at the same time.

WHAT IS A BUFFER? A buffer is a solution of a
weak acid or base and its salt [which is its
conjugate]

WHAT DOES IT DO? A buffer resists a change in pH

HOW DOES IT WORK? Since a buffer consists of
both an acid or base and its conjugate, an acid
and a base are present in all buffer solutions. If a
small amount of strong acid is added to the
buffer, there is a base component ready and
waiting to neutralize the “invader”

PREPARING BUFFER SOLUTIONS:

Use 0.1M to 1.0M solutions of reagents and
choose an acid whose Ka is near the [H+] or
[H3O+] concentration we want. The pKa should
be as close to the pH desired as possible. Adjust
the ratio of weak A/B and its salt to fine tune the
pH.