Topic 9: Acids and Bases
9.1 Properties of Acids and Bases
9.1.1. Outline the characteristic properties of acids and bases in aqueous solution.
The properties that must be considered are: effects on indicators and reactions of acids
with bases, metals and carbonates. Bases which are not hydroxides, such as ammonia,
soluble carbonates and hydrogencarbonates should be included. Alkalis are bases that
dissolve in water.
It is kind of hard to describe acids and bases without using either Bronsted-Lowry or
Arrhenius, but I guess we will just describe what they do instead of what they are.
Indicators are special substances that are added to solutions to tell what pH (is it acidic or
basic?) that solution is. You don’t need to know any specifics right now, just know that
indicators change colors depending on the pH of the substance they are being mixed with.
When acids mix with bases they USUALLY (not ALWAYS) produce water (HCl +
NaOH  H2O + NaCl). When acids mix with metal they normally produce hydrogen
gas (2HCl + Mg  MgCl2 + H2). When acids mix with carbonates they produce water
and CO2 (2HCl + CaCO3  CO2 + H2O + CaCl2). The Carbon Dioxide is the fizzling
that you see when you mix an acid with baking soda, carbon dioxide is being released in
the form of a gas.
Some examples of acids are: HCl (hydrochloric acid), CH3COOH (ethanoic acid), H2SO4
(sulfuric acid), NH4+ (ammonium). Some examples of bases are NaOH (sodium
hydroxide), NH3(ammonia), and CH3COO-.
The term alkaline applies to bases dissolved in water (an alkaline mixture is a basic
mixture).
9.2 Strong and Weak Acids and Bases
Note: Bronsted-Lowry definitions of acids and bases are not required for this subtopic.
9.2.1. Describe and explain the differences between strong and weak acids and bases in
terms of the extent of dissociation, reaction with water and conductivity.
The term ionization can be used instead of dissociation. Solutions of equal concentration
can be compared by pH and/or conductivity.
Strong and weak acids are defined by their ease of losing (or donating) a proton. A
strong acid, when placed in water, will almost fully ionize/dissociate straight away,
producing H3O+ ions from water. A weak acid will, however, only partially do this,
leaving some unreacted acid remaining. Strong acids are much more conductive then
weak acids because they dissociate completely and leave a lot of ions in water, which
make them conducive while weak acids have much less ions in water and thus much less
conductivity. Strong bases will accept all of the hydrogen ions in the reaction until they
are completely used up, while weak bases will only absorb some of them. A strong base
will rip all the hydrogen it can from water molecules, leaving OH- behind, while weak
bases won’t do it as much. Strong bases also create more ions in solution, and are thus
more conductive.
9.2.2. State whether a given acid or base is strong or weak.
Specific strong acids are hydrochloric acid, nitric acid and sulfuric acid. Specified weak
acids are ethanoic acid and carbonic acid (aqueous carbon dioxide.)
Specified strong bases are all group 1 hydroxides and barium hydroxide. Specified weak
bases are ammonia and ethylamine.
Strong Acids are: HCl, HNO3, H2SO4.
Weak Acids are: CH3COOH, H2CO3.
Strong Bases: Any group 1 hydroxide (ie NaOH, etc.), BaOH.
Weak Bases: NH3, CH3CH2NH2.
9.2.3. Describe and explain data from experiments to distinguish between strong and
weak acids and bases, and to determine the relative acidities and basicities of
substances.
One way to observe the strength of an acid or base is to observe its reactions where you
know what the products will be. If an acid very quickly produces carbon dioxide when
mixed with carbonates, then it is pretty strong, but if it is much weaker then it is not so
strong. Same with metals and producing hydrogen gas. Other ways you can tell are by
using pH paper or a meter, which tells you how strong/weak and acid or base is by giving
you a value on the pH scale. The relative acidities (I’m assuming that means diprotic or
something) can also be found by neutralizing two acids with a strong base in the presence
of an indicator.
9.3 The pH scale
9.3.1. Distinguish between aqueous solutions that are acidic, neutral or basic using the
pH scale.
On the pH scale, acidic aqueous solutions will have a value of less then 7. Neutral
aqueous solutions will have a value of 7. Basic aqueous solutions will have a value that
is greater then 7.
9.3.2. Identify which of two or more aqueous solutions is more acidic or basic, using
pH values.
Measure pH using a pH meter or pH paper. Students should know that pH paper
contains a mixture of indicators. The theory of pH meters is not required.
So, you have two substances and you want to know which one is more acidic or basic.
How you normally do this is by using either pH paper or a pH meter. pH paper is a piece
of paper that has in a mixture of indicators. When you insert it into a certain substance,
the pH of that substance will turn one of the indicators in pH paper a certain color and
you will observe that color and know that the substance if of a certain pH by comparing it
with a scale that is brought with pH paper. You could also use a pH meter and it will tell
you exactly the pH of the solution. The solutions with more extreme values (lower or
higher) are the ones that are more acidic or basic respectively.
9.3.3. State that each change of one pH unit represents a tenfold change in the
hydrogen ion concentrations [H+(aq)].
Relate integral values of pH to [H+(aq)] expressed as powers of ten. Calculation of pH
from [H+(aq)] is not required.
The pH scale is a log scale, so each change of one pH unit represents a tenfold change in
the hydrogen ion concentrations in the solution [H+(aq)](this value is how we get pH, more
in HL).
9.3.4. Deduce changes in [H+(aq)] when the pH of a solution changes by more than one
pH unit.
So let’s say your pH value changes by 2, from 5 to 3. At 5, you have .00001 moles per
liter of hydrogen ions. At 3, you have .001 moles per liter of hydrogen ions. Let’s say
you go from 5 to 3.5. At 5 you have .00001 moles per liter of hydrogen ions.
Concentration is 10(-pH) moles/liter.
9.4 Buffer Solutions
9.4.1. Describe a buffer solution in terms of its composition and behavior.
A buffer resists change in pH when a small amount of a strong acid or base is added.
Suitable examples include ammonium chloride/ammonia solution and ethanoic
acid/sodium ethanoate. Blood is an example of a buffer solution.
A buffered solution is one that resists a change in its pH when either hydroxide ions or
protons are added. A buffered solution may contain a weak acid and its salt (for example,
CH3COOH and NaCH3COO) or a weak base and its salt (for example, NH3 and NH4Cl).
A good example of a buffered solution is your blood. It resists change in pH in either a
more acidic or basic direction and if it didn’t, your cells and proteins would stop working
and you would die very quickly.
9.4.2. Describe ways of preparing buffer solutions.
9.5 Acid-base titrations
9.5.1. Draw and explain a graph showing pH against volume of titrant for titrations
involving strong acids and bases.
Graphs below. Notice that on both graphs the curve starts off going decreasing or
increasing fairly slowly until it gets close to the equivalence point. This is because when
the titration first begins there is a lot of H+ or OH- ions in solution already, and the titrant
doesn’t have much effect. However as you near the equivalence point the amount of ions
that were already in there begins to drastically decrease, and then for a little bit the ions
you add change the pH drastically because they are the only ions in solution, but then
when you keep adding the more they have less of an effect because there is already a lot
in them and it takes a considerable amount of ions to increase the pH (remember it is on a
log scale, to change pH one number at the more extreme ends requires a LOT more ions
then it takes to change it at the equivalence point.)