Net Ionic Equation
• Net ionic equations are used to show only the
chemicals and ions involved in a chemical
reaction in order to simplify information about
a reaction.
• The ions that are not involved in the reaction
are called spectator ions and are removed
from the reaction.
NaOH(aq) + HCl(aq)  NaCl(aq) + H2O(l)
+
+
+
Na + OH + H +Cl  Na +Cl + H2O(l)
OH- + H +  H2O(l)
Acids
• Arrhenius definiton of acids
– A substance that produces H+ ions
• Brønsted-Lowry definition of acids
– An acid is a substance that can transfer a hydrogen
ion (H+), also called a proton, or Hydronium (H3O+)
to another substance. It is often called a “proton
donor.”
Notice the slight difference in wording, but
essentially they mean the same thing…
Ex: HCl  H+ + ClHNO3  H+ + NO3HCH3CO2  H+ + CH3CO2-
The Hydronium Ion
• H3O+ is called the hydronium ion.
• It is another way to understand how hydrogen
+
ions, H , move around in a solution.
• Hydronium ions are created when an acid is
dissolved into water.
HCl(g) + H2O(l)  H3O+(aq) + Cl-(aq)
HNO3 (s) + H2O(l)  H3O+(aq) + NO3-(aq)
Naming Acids
• Binary acids – usually acids that are made of
hydrogen and one other substance.
• Name these by using hydro- for the hydrogen and
the name of the second element with ending of –ic,
then following it up with the word, acid.
Ex: HCl – hydrochloric acid
Usually binary means 2, but exceptions are made
when it is an acid of a polyatomic ion that does not
contain oxygen
Ex: HCN – hydrocyanic acid
Naming Acids
• Oxyacids – acids of polyatomic ions that contain
oxygen.
• Name these by noting the name of the oxyanion and
following it with the word acid.
If it ends in –ate, then the acid is named by the
oxyanion and the suffix of –ic.
If it ends in –ite, then the acid is named by the
oxyanion and the suffix of –ous.
Ex: HBrO3 – bromic acid
HBrO2 – bromous acid
What we know about acids
• Acids are well known as substances that are
corrosive, sour tasting, turn litmus paper a red color,
and react with bases and metals to form salts.
• Common acids that you deal with every day include:
Acetic acid, carbonic acid, citric acid, maleic
acid, fumaric acid, ascorbic acid, acetylsalicylic
acid, amino acids, hydrochloric acid
Bases
• Arrhenius definition of bases
– A substance that produces OH- ions
• Brønsted-Lowry definition of bases:
– a substance that can accept a hydrogen ion (H+)
from a proton donor. It is often called a “proton
acceptor.”
Notice that these two definitions are very different.
Ex:
NaOH + H+  Na+ + H2O
NH3 + H+  NH4+
Mg(OH)2 + 2H+  Mg+ + 2H2O
NaHCO3 + H+  Na+ + H2CO3
What we know about bases
• Bases are commonly known as substances that are corrosive,
have a slippery feel, taste bitter, turn litmus paper blue, and
react with acids to form salts.
• Basic solutions are often called “alkaline solutions.”
• Common bases that you may have dealt with include:
Ammonia, sodium bicarbonate, sodium
hydroxide, magnesium hydroxide, calcium
hydroxide, calcium carbonate
(coffee is acidic… you’ll notice this when you drink bad
coffee)
Strong vs Weak
• Acid and base strengths are determined by their
abilities to dissociate (ions breaking apart when
dissolved)
• Strong acids and bases completely dissociate. Not
only are these powerful chemicals, but are also
extremely good electrolytes (ie: they conduct
electricity extremely well)
• Weak acids and bases do not completely dissociate
meaning that fewer ions enter into solution. A
noticeable feature of weak acids and bases are that
they are poor electrolytes (ie: they do not conduct
electricity well)
pH
• pH is defined as the measurement of the hydrogen ion
concentration present in a solution.
• Mathematically, it stands for:
pH = -log [H+]
• The typical pH scale ranges from 0 to 14
• 0 - 6.9 Acidic range. (0 = strongest)
• 7
Neutral (not acidic or basic)
• 7.1 – 14 Basic range. (14 = strongest base)
• The pH of some common materials:
ACIDS
HCl
Vinegar
Lemon Juice
Apple juice
Acid rain
Human saliva
-1 to 0.1
2.2
2.2 to 2.4
2.9 to 3.3
5.2
6.3 to 6.6
BASES
Human blood
Sea water
Eggs
Baking soda
Borax
Soda lye
7.35 to 7.45
7.36 to 8.21
7.6 to 8
8.0
9.2
13.0 to 15.0
Water has neutral pH
• Water has a natural ability to dissociate into H3O+
and OH- ions.
H2O + H2O  H3O+ + OH• Kw, the self-ionization constant of water, says there
will always be a natural balance of H3O+ to OH- ions
in water.
Kw = [OH-] x [H3O+] = 1.00 x 10-14
• Pure water has 1.00 x 10-7M OH- and 1.00 x 10-7M
H3O+
(do the math… why is the pH scale 0-14? Why is 7
neutral?)
Acid Dissociation Constant
• Strength of acids are determined by numerical
values called the acid dissociation constant
(Ka)
• This basically tells you how well an acid will
dissociate in water
• The bigger the number, the stronger the acid
• The smaller the number, the weaker the acid
Acid and Conjugate Base
• We can show the dissociation of an acid and write it
as an ionic reaction
H2S  H+ + HS- Ka = 8.9x10-8
• Notice that in this equation, H2S donates a H+ ion, so
is thus defined as an acid.
• The reverse reaction can also occur
Kw
H+ + HS-  H2S Kb =
8.9 x10 8
• Now notice in this next reaction, HS- accepts a H+ ion
so it fits the definition of a base. We can say that HSis the conjugate base of H2S
Base Dissociation Constant
• Similarly, base strength is characterized by
how well bases will gain H+ ions.
• This is characterized numerically by the base
dissociation constant Kb
• The larger the value of Kb, the stronger the
base
• The smaller the value of Kb, the weaker the
base.
Base and Conjugate Acid
• We can show the dissociation of a base and write it
as an ionic reaction
NH3 + H2O  NH4+ + OH-
Kb = 2.5x10-5
• Notice that in this equation, NH3 accepts a H+ ion, so
is thus defined as a base.
• The reverse reaction can also occur
Kw
NH4+ + OH-  NH3 + H2O
Ka =
2.5 x10 5
Now notice in this next reaction, NH4+ donates a H+
ion so it fits the definition of an acid. We can say
that NH4+ is the conjugate acid of NH3
Practice
• Balance the following equation:
NaOH(aq) + H2SO4(aq)  Na2SO4 (aq) + H2O(l)
• What are the acid, base, conjugate acid,
conjugate base of the above reaction?
• Write the net ionic equation for the above
reaction
• How does this net ionic equation exemplify
the reaction of acids with bases?
NaOH + H2SO4Na2SO4 + H2O
How do you identify the acid in a chemical equation?
•
Look at the reactants
•
Which reactant has Hydrogen?
•
Which reactant has Hydrogens to be given away?
NaOH
or
H2SO4
How do you identify the base in a chemical equation?
•
Look at the reactants
•
Which reactant has hydroxide?
•
Which reactant may want to absorb hydrogen?
NaOH
or
H2SO4
•
•
•
•
•
•
How do you identify the conjugate acid in a chemical
equation?
Look at the products.
Which product has Hydrogen?
Which product has Hydrogens that can be removed to
recreate the original base?
Na2SO4
or
H2O
Look at the products.
Which product has looks like it’s missing Hydrogens?
Which product will recreate the original acid if we give
hydrogens back to it?
Na2SO4
or
H2O
Neutralization
• Neutralization occurs when an acid reacts with a base
to produce a salt and water.
• A complete neutralization occurs when all the available
H+ ions from the acid are reacted, and no excess base is
present. In other words, the amount of acid = the
amount of base in the reaction solution. This is also
known as the equivalence point
• An acid-base indicator is a substance whose color is
affected by acidic and basic solutions. We can use
indicators to monitor pH changes in an acid-base
reaction.
Titration
• Titration is a technique that uses acid-base
neutralization to determine the amount of a
substance in solution as it reacts with a known
amount of acid or base.
• Ex: I have a 10mL solution of NaOH of
unknown concentration. It reacts completely
with 35mL of 1.00M HCl solution. What is the
molarity of the NaOH solution?
Monoprotic vs Polyprotic
• Many acids are simple combinations of one
H+ ion and one -1 ion. With these, the acid
can donate no more than one H+ to another
substance. These are called monoprotic acids
• Other acids have more than one H+ ion
bonded to the negative ion (meaning -2 or -3
anion). These can give up one H+ or more
than one H+ ions to another substance. These
are called polyprotic acids.
Polyprotic acid dissociation
• Dissociation – process by which ionic
compounds separate into smaller ions
• Step 1
H3PO4 + H2O  H2PO4- + H3O+
• Step 2
H2PO4- + H2O  HPO42- + H3O+
• Step 3
HPO42- + H2O  PO43- + H3O+
complete dissociation
• Net Reaction
H3PO4 + 3H2O  PO43- + 3H3O+
Acid-Base Reactions
• When acids react with metals, metal salts and
hydrogen gas are formed.
• 2HCl + Mg  MgCl2 + H2
• H2 SO4 + 2Na  Na2 SO4 + H2
– In general: HX + M  MX + H2
• When acids and bases react, we simply observe a
transfer of a H+ ion
• HNO3 + NaOH  NaNO3 + H2O
• HCH3CO2 + NaHCO3  NaCH3CO2 + H2CO3
• When we see a substance that can potentially act
as an acid AND a base, it is called amphoteric.
• H2O + H2O  H3O+ + OH-