Chemistry I Honors
Acids and Bases
Lesson #1
Introduction
• Welcome to Acids and
Bases !
• This unit will use a lot of
the concepts covered in
the Solutions unit, but
also contains some very
new material.
• We will look at the
characteristics and
behaviors, and reactions
of both types of
compounds.
First – Some History and Vocabulary
• In the mid-1800’s ,
scientists discovered that
some compounds would
dissolve in water and the
resulting solution would
conduct electricity.
• Compounds that behaved
this way were classified as
electrolytes.
• Ionic compounds (like
salt as NaCl, potassium
nitrate as K(NO3), etc)
were found to be
electrolytes.
• It was also taken as
“common knowledge”
that covalent compounds
did not produce solutions
that conducted electricity
– they were nonelectrolytes.
How Does This Occur?
• Even though science
could identify which
compounds were
electrolytes, there was
no understanding about
what happens when
one of these
compounds dissolves
that could account for
the electrical
conductivity.
A Hypothesis
• In the late 1800’s, a Ph.D. candidate from
Denmark proposed that ionic compounds break
apart to form charged particles when dissolved in
water.
• He named the process “Electrolytic Dissociation”.
• The idea was not well received – it was believed
at the time that compounds were built from
neutral atoms – and the electron had not yet
been discovered by Thomson.
Eventually…
• Sub-atomic particles are
discovered, and the
theory of Electrolytic
Dissociation is
accepted.
• In the early 1900’s the
Nobel Prize in
Chemistry is awarded to
Svante Arrhenius for his
theory of electrolytic
dissociation.
What This Means
• Ionic compounds like NaCl are electrolytes
because they dissociate in water to form
charged ions.
NaCl (s)  Na+1 (aq) + Cl-1 (aq)
• Covalent compounds like sugar are nonelectrolytes because they are molecular –
their dissolving yields complete molecules
that do not have a charge.
C6H12O6 (s)  C6H12O6 (aq)
What Does This Have to do with Acids?
• Acids are covalent
compounds.
• We use ionic strategy to
write their formulas
(SOCCR), but they are
covalent compounds
(technically polar
covalent).
• But (and this is huge),
they are electrolytes.
• Solutions of acids
conduct electricity – as
if they are ionic – but
they aren’t…
• And, since they are
covalent, they do not
dissociate into ions like
ionic compounds do.
Where do the Ions Come from in
Acids?
• If acid solutions conduct electricity, something
has to occur to produce ions in solution.
• Arrhenius proposes (later found correct) that
Acids actually react with the water to
produce ions.
• This reaction is named “Ionization” – the
action of forming ions.
The Reaction
• We have learned that Acids have the general
formula “HX” , where the “X” part is either a nonmetal element or a polyatomic ion.
• From bonding, the H and the X share a pair of
electrons in a covalent bond.
• During the Ionization reaction, the H – X bond is
broken and the H becomes attached to the two
lone pairs on the Oxygen atom of the water
molecule.
• But, (and this is critical) the X atom keeps both
electrons from the H – X bond.
Hydronium
• The H atom leaves his electron behind (with
the X atom), meaning he now has a +1 charge.
• The X atom has gained the H’s electron,
meaning that he now has an extra electron
and as a result now has a -1 charge.
• The H+1 particle connects to the water
molecule, forming an ion that has the formula
H3O+1
• This ion is called the “hydronium ion”.
Equation
• In equation form, the ionization of an acid HX
will look like this:
HX + H2O  H3O+1 (aq) + X-1 (aq)
• All acids that have the chemical format “HX”
will react the same way – will all form
hydronium and an anion.
• Example – Hydrochloric Acid
HCl + H2O  H3O+1 (aq) + Cl-1 (aq)
A Visual
More Examples
• For Hydrobromic Acid:
HBr + H2O  H3O+1 (aq) + Br-1 (aq)
It works for oxy-acids too!
• For Nitric Acid:
H(NO3) + H2O  H3O+1 (aq) + (NO3)-1 (aq)
Some Quick Vocabulary
• Monoprotic – this is an
acid that only has 1 H
atom bonded to the “X”
part of the acid
molecule.
• Examples would be
–
–
–
–
HCl
HBr
H(NO3)
H(C2H3O2)
• Diprotic – these are
acids that have 2 H
atoms bonded to the
“X” part of the acid
molecule.
• Examples would be
– H2(SO4)
– H2(CO3)
Diprotic Ionizations
• Important to remember that during any
ionization, the acid loses 1 H atom to a water,
forming a single hydronium ion.
• This holds true for diprotic acids, but they are
able to undergo two separate ionizations.
#1) H2(SO4) + H2O  H3O+1 + H(SO4)-1
#2) H(SO4)-1 + H2O  H3O+1 + (SO4)-2
• Look carefully through the sequence presented
above. See if you can write all three ionization
equations for phosphoric acid: H3(PO4)
What About Bases?
• Bases are easier to deal
with.
• Arrhenius defined bases
as “Metal Hydroxides”
that dissociate in water to
produce a metal cation
and the hydroxide anion.
• The dissociation is exactly
the same as we learned in
the Solutions Lesson #1.
• There are “monobasic”
compounds that have one
hydroxide attached to the
metal cation.
• And there are “dibasic”
compounds that have two
hydroxides attached to
the metal cation.
Monobasic and Dibasic Dissociations
• Na(OH) is monobasic. (got the formula from
SOCCR)
• It’s dissociation looks like this:
Na(OH) (s)  Na+1 (aq) + (OH)-1 (aq)
• Mg(OH)2 is dibasic. (again, formula from SOCCR)
• It’s dissociation looks like this:
Mg(OH)2 (s)  Mg+2 (aq) + 2 (OH)-1 (aq)
Summary
• Acids:
– Are electrolytes
– Have the formula HX
– Ionize in water to form
hydronium
– Can be monoprotic,
diprotic, or triprotic
• Bases:
– Are electrolytes
– Have the formula M(OH)
– Dissociate in water to for
aqueous ions – most
importantly the OH-1 ion.
– Can be monobasic or
dibasic