Aqueous chemistry is a very important component to laboratory activity. Liquid solutions are
more convenient to keep and mix together than either gases or solids. A typical solution
consists of a smaller amount of one substance, the solute, dissolved in a larger amount of
another substance, called the solvent. When the solute dissolves in the solvent, its entities
become evenly dispersed throughout the liquid. The concentration of a solution is usually
given in the amount of solute dissolved in a given amount of solution.
Molarity: moles of solute dissolved in Liters of solution
M = moles solute
L solution
Concept Tests:
What is the concentration, in moles per L of 45 grams of hydrobromic acid dissolved in 300 ml
of water?
How many grams of solute are in 2.25 L of 0.456 M potassium chloride?
Stoichiometry for solution chemistry can be used. The link between volume and moles is the
Molarity of the solution. When given mL, or L, you can directly get moles by multiplying by
the concentration (in moles/L). Just make sure your volumes cancel out!!! (convert mL  L)
Water is known as the universal solvent for a reason. It is a polar molecule which has the ability
to act as a solvent for polar compounds as well as ionic compounds. It even has the ability to
dissolve some nonpolar gases!
How ions in solution behave with respect to water molecules:
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The partially negative oxygen atom in water will surround the positive cations, while the
partially positive hydrogen atoms will surround the negative anions. If two electrodes
(conductors of electricity) are placed in a solution and connected to a battery, the cations will
migrate through the solution to the negatively charged electrode and the anions will migrate
towards the positively charged electrode. If a light bulb is inserted into the circuit, the bulb will
light up, showing that the solution can conduct electricity. Compounds whose aqueous
solutions conduct electricity are called electrolytes. All ionic compounds that are soluble in
water are electrolytes.
Electrolytic solutions can be taken a step further. Is the substance that was dissolved (the
solute) a good electrolyte or a poor electrolyte? If something is a good electrolyte is completely
or almost completely dissolves in solution to create ions. Thus, it will then be a good conductor
of electricity. Good electrolytes are called strong electrolytes. If a solute does not dissolve well
in solution, then there will not be many ions floating around, and as such, it will not conduct
electricity well. Poor electrolytes are called weak electrolytes. Finally, there are substances that
do dissolve in water (completely even!!) but still do not conduct electricity. The reason is, there
are no ions formed in solution to carry the electrical current. Species that do dissolve in solution
but do not conduct electricity are called nonelectrolytes.
Strong Electrolytes
Ammonium hydroxide
Sodium chloride
Strong Electrolyte
Weak Electrolytes
Acetic acid
Barium hydroxide
Weak Electrolyte
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Nonelectrolytes
Sugar
Ethanol
Nonelectrolyte
Of all the chemical reactions that occur in nature, the majority of them take place in aqueous
media (water) and many of them involve ions. Chemists use three types of equations, in a stepby step process to indicate the chemical reaction that occurs. The first is the molecular: or the
overall reaction. The second is the ionic, where we focus on what is actually reacting (or not
reacting!) and the third is the net ionic, where we ignore species that are not actually involved in
the chemical reaction and simply indicate the species that do chemically change.
It is important to remember what it means to put a salt in water. The salt solid dissociates into
IONS. These ions are now independent of one another and free to move and tumble about the
water solution.
In precipitation reactions, two soluble ionic compounds come together to form a solid, insoluble
product, called the precipitate. Precipitates form for the same reason that some ionic
compounds do not dissolve in water: the electrostatic attraction between the ions outweighs the
tendency of the ions to become solvated and interact in an ion-dipole manner with water.
When solutions of such ions are mixed, the ions collide and stay together, and the resulting
substance is said to “come out of solution”.
In order to determine if a precipitate is going to form, you need to examine the equation. Are
the original compounds soluble in water? What if one reactant was NOT soluble in water?
What would that mean in terms of a chemical reaction taking place? If both reactants are
soluble in water, you must examine the ions that are formed in solution. Then examine all
possible combinations of ions. This is the double replacement idea. The cation-anion partners
will “switcharoo”. In order to determine if a possible ion combination is soluble in water you
will have to be aware of the solubility rules. If you form an ion combination that is NOT soluble
based off the solubility rules then you will have a net-ionic equation. Otherwise, you will only
keep all ions in solution (as seen in the ionic equation).
Solubility:
What compounds are soluble in water? Here is a nice little list for you.
1. All common compounds of Group I and ammonium ions (NH4+1) are soluble.
2. All nitrates (NO3-1), acetates (C2H3O2-1) , and chlorates (ClO3-1) are soluble.
3. All binary compounds of the halogens Group VIIA (other than F) with metals are soluble,
except those of Ag, Hg(I), and Pb. Pb halides are soluble in hot water.
4. All sulfates are soluble, except those of barium, strontium, calcium, lead, silver, and mercury
(I). The latter three are slightly soluble.
5. Except for rule 1, carbonates (CO3-2), hydroxides (OH-1), oxides (O-2), silicates, and
phosphates (PO4-3) are insoluble. Calcium, strontium and barium hydroxide are strong
bases and therefore are soluble ENOUGH to be aq in solution
6. Sulfides (S-2) are insoluble except for calcium, barium, strontium, magnesium, sodium,
potassium, and ammonium.
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For a double replacement reaction, the ions switch partners. Whatever anion the first positive
ion was bonded to switches and bonds with the second anion. The second cation switches and
bonds with the first cation’s anion. When writing the molecular equation we want to indicate
the species that are coming together and “doing” the reacting.
Concept Check:
Write the potential products in this double replacement reaction. Do not forget to
balance the equation!
AgNO3 (aq) + Na2CO3 (aq)  ?????? +
??????
This overall reaction is the molecular equation. But what is really going on in this aqueous
state? Notice, we have formed a solid. That means that ions that were once dissolved in
solution came together and made a solid, or a precipitate. The ionic equation will help us
determine what is going on in solution. Since our two reactants are aqueous that means that
they are dissolved in water. We can therefore write them as individual ions. Be careful of the
charges, the number of each, and do not forget to indicate their phase!!!!
Remember, ionic means IONS, the compound silver carbonate is NOT soluble in water, it is a
solid, that means we cannot split it apart into its ions! From this ionic equation we can see that
some ions appear on both sides of the arrow (the reactant and the product side). Well, they just
do not seem to do all that much now do they? It is almost as if they just “watch” the chemical
reaction that takes place and forms the silver chromate solid. These ions are termed spectator
ions for that very reason. Any ion that appears in the same state on the left side (reactant side)
and the right side (product side) of the arrow can, in effect, be canceled out of the equation.
That is exactly what you do when you write the net-ionic equation. Get rid of the spectator ions
and write the ions that actually “do” something.
The net-ionic equation is most useful because it eliminates the spectator ions (the chemical
species that are not doing anything) and shows the actual chemical change taking place. In fact,
if we changed our reagents to be K2CrO4 (aq) and AgC2H3O2 (aq) (potassium chromate and silver
acetate) the overall chemical reaction would stay exactly the same, only the spectator ions have
changed. Take a moment and write out the chemical equation, balance it, write the ionic and
net-ionic equations to prove this fact.
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Acids and bases are important compounds. We have talked about how to name acids
specifically the binary halogen acids (HF, HCl, HBr, and HI as being names hydrofluoric acid,
hydrochloric acid, hydrobromic acid, and hydroiodic acid) and the oxoacids (H2SO4, H2CO3,
HNO3, and HNO2 for example: sulfuric acid, carbonic acid, nitric acid, and nitrous acid – just to
name a few!).
In 1884 Svante Arrhenius proposed the first theoretical model for acids and bases. Prior to that
time, these chemically opposite substances were described in properties such as their taste; their
effects on metals, carbonates, and dyes (called indicators); their feel to the touch, and their
ability to react with each other. According to the Arrhenius theory, pure water dissociates to
some extent to produce hydrogen ions, H+ and hydroxide ions, OH-. When this occurs, equal
amounts of H+ and OH- ions are produced:
H2O(l)  H+(aq) + OH-(aq)
An acid, according to Arrhenius, is any substance that liberates H+ ions when placed in water.
When the H+ concentration is elevated this solution is said to be acidic. Similarly, a base is
defined as any substance that liberates OH- ions when placed in water. The resulting solution
has a higher concentration of OH- ions than H+ ions and is said to be basic, or alkaline.
The strength of an acid or base is a measurement of the extent to which reactions such as the
two above react. A strong acid or base is said to dissociate 100% (and is thus called a strong
electrolyte!). For example HCl, hydrochloric acid dissociates fully into H+ and Cl- ions. There
are no HCl molecules left, they have all separated. The same goes for strong bases like the
sodium hydroxide, NaOH there will not be any NaOH left it will all be dissolved into Na+ and
OH- ions.
HCl(aq) + H2O(l)  H+(aq) + Cl –(aq)
NaOH(aq)  Na+(aq) + OH-(aq)
H
H
H N
H
O
H
H N H
-
OH
H
H
Again, the reaction of HCl takes place almost completely. So, a 2.0 Molar HCl solution
produces 2.0 M [H+1] and 2.0 {M] Cl-1 ions when dissolved in water. Thus, hydrochloric acid is
classified as a strong acid. On the other hand, the reaction of ammonia, NH 3, to form the NH4+
ion does not take place completely. Most of the ammonia remains in the form of NH 3
molecules. Thus, 2.0 Molar NH3 dissolved in water produces less than 2.0 M [NH4+1] and less
than 2.0 M [OH-1] ions. Because of this, ammonia is called a weak base (and thus, is called a
weak electrolyte!).
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The list of strong acids include, HCl, HBr, HI, HNO3, H2SO4 and HClO4. The list of strong bases
include, LiOH, NaOH, KOH, RbOH and CsOH. What do the strong bases have in common?
They are all composed of a group one metal and a hydroxide. From the solubility rules all
compounds containing a group one metal are soluble. All of these acids and bases dissociate
100% in water.
Common Strong Acids and Strong Bases
ACIDS
Group VII hydrides
Oxoacids
HCl
HNO3
HBr
H2SO4
HI
HClO4
BASES
Group IA hydroxides
LiOH
NaOH
KOH
RbOH
CsOH
Group IIA hydroxides
Mg(OH)2
Ca(OH)2
Sr(OH)2
Ba(OH)2
The term neutralization is often used to describe a reaction in which equal amounts of acid and
base react with each other. Acid and bases can react with some compounds to change their
color, these compounds are dyes that change color as the pH changes are referred to as
indicators. This is one method used to determine the point at which exact amounts of acid and
base have been reacted in neutralization reactions.
Normally, the products of the reaction are water and salt. The salt follows the rule of solubility
and will either dissolve or not depending on what salt if formed. Thus, the neutralization of
nitric acid with potassium hydroxide is represented by the following reaction:
HNO3(aq) + KOH(aq)  H2O(l) + KNO3(aq)
The essential change is all aqueous reactions between strong acids and bases is that the H+1 ion
from the acid, and the OH-1 ion from the base form a water molecule. Thus, for all
neutralization reactions, the net-ionic equation looks the same:
H+1(aq) + OH-1(aq) → H2O(l)
Concept Test:
Write the ionic and net ionic equations for the above acid-base reaction:
H+1(aq) + NO3-1(aq) + K+1(aq) + OH-1(aq)  H2O(l) + K+1(aq) + NO3-1(aq)
H+1(aq) + OH-1(aq)  H2O(l)
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In a typical aqueous neutralization reaction the products are a salt and water. If the water was
evaporated, we would be left with KNO3 as a solid salt.
In a titration, a solution of a known concentration is used to determine the concentration of
another solution through a monitored reaction. In an acid-base titration, a standardized (known
concentration) of base is added slowly, and in known volumes to an acid solution of unknown
concentration.
In an acid-base titration, an indicator is used to show the change from an acidic situation (all
acid), and as the base is slowly added and neutralization occurs, the color shift will be towards
the basic side. When the moles of acid = moles of base neutralization is said to occur. We note
this by the change in the indicator. Without the indicator, the solutions might be colorless and
clear, which would make determining the neutralization point very difficult!
The equivalence point is when the moles H+ = moles OHThe endpoint occurs when we have a tiny excess of base in the solution which has caused our
indicator to change color. The amount of base needed to reach the endpoint is the same amount
of base needed to reach the equivalence point. You can use the known information of the
concentration of the base, the amount of base added (in mL) to reach the equivalence point, and
the fact that at the equivalence point, moles OH-1 = moles H+1 in order to determine the
concentration of the acid. But again, you must have a balanced chemical equation!!
In the titration above both the acid and the base are clear and colorless, addition of the base (in the buret), to the
acid (in the Erlenmeyer), shows a color change occurring with the indicator. The indicator is colorless in acid, but
in the presence of base, turns pink. When doing a titration, your goal is to have the moles H + = moles OH-.
Therefore, ideally the solution should be the LIGHTEST pink possible indicating the smallest excess of OH - present.
This is the endpoint of the reaction and shows the equivalence point, where the moles of H + = moles OH-.
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General Steps to Solve Acid-Base Problems:
1. Write the balanced chemical equation
2. Determine the volume of base/acid added
3. Using the molarity and volume, determine the moles of acid or base
4. Since moles OH-1 = moles H+1 determine the moles at the equivalence point using the
mole:mole ratio
5. Knowing the volume and the moles, determine its molarity
Sometimes, when species are mixed together, a gas evolves. How many of you have taken
Alka-seltzer? When you open the package there are two tablets there. They are not bubbling.
They are not doing much of anything – but when you drop those tablets in water, then away go
the gas bubbles.
The active ingredients in Alka-Seltzer are sodium bicarbonate, NaHCO3, and citric acid,
H3C6H5O7. When water is added to an Alka-Seltzer tablet, the bicarbonate ion acts as a base to
accept the two hydrogen ions from citric acid and form carbonic acid and the citrate ion. The
carbonic acid immediately breaks down to form carbon dioxide gas and water. When taken for
an upset stomach, the citrate ion acts as a base to neutralize the excess stomach acid.
3 HCO3-1 (aq) + H3C6H5O7 (aq) → 3
CO2 (g) + 3 H2O(l) + C6H5O7-3(aq)
What about when you make bread, or a cake, or muffins? When you put the dough or batter
into the oven it is flat. But after some cooking, you removed your food item and it has “puffed”
up. Why did it puff? Gas evolved! Several types of reaction lead to gas formation. The most
common gas formed it CO2 but H2 and CO can also be formed.
Any time a metal carbonate (M+x CO3-2) is placed in an acidic solution, a gas is evolved. This is
due to the fact that the acid can dissolve the metal carbonate (if it is not already dissolved!) and
formed H2CO3 (aq). In a short amount of time (fairly instantaneously on our time scale) the
carbonic acid decomposes into water and carbon dioxide gas. The formation of the carbon
dioxide gas is the NET-IONIC equation for this type of reaction.
Gas Formation:
Molecular: CaCO3 (s) + 2HCl(aq) → CaCl2 (aq) + H2CO3 (aq)
Net-ionic: H2CO3 (aq) → H2O(l) + CO2(g)
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Redox reactions, otherwise known as oxidation-reduction reactions are another type of chemical
reaction to be aware of. These type of reactions include the formation of a compound from its
elements (and vice-versa), all combustion reactions, reactions that generate electricity in
batteries, reactions that produce cellular energy, and many others.
In redox reactions, the key chemical event is the movement of electrons from one species to
another. The driving force for the movement of electrons occurs from the reactant with the
lower attraction for the electron(s) to the reactant with more attraction for the electron(s).
Examine the following reaction:
Mg(s) + O2 (g) → 2MgO(s)
Magnesium is a neutral element/atom. O2 is also a neutral molecule. But when these two
neutral species come together in a chemical reaction, they form an ionic compound. On the
product side, we have formed a charged Mg+2 ion which is attracted to the O-2 ion. This means
that Mg metal as the reactant must have given up (lost) its electrons to the oxygen molecule
reactant.
There is some terminology for redox reactions that you should be aware of – and it can be quite
tricky!!
Oxidation is the loss of electrons. When something loses electrons is becomes positive, or more
positive as the case may be.
Reduction is the gain of electrons. When something gains electrons it becomes negative, or less
positive, whatever the case may be.
Here is the confusing part. Since oxidation can only occur with reduction, and reduction must
occur with oxidation the two are linked together in the chemical reaction. If one species is
oxidized, another must be reduced. Thus, the species that is oxidized is often termed the
reducing agent, and the species that is reduced is called the oxidizing agent.
Mg(s) + O2 (g) → 2MgO(s)
Oxidation: Mg(s) → 2Mg+2 + 4e-1
Reduction:
O2 + 4e-1 → 2O-2 (!!)
Since oxygen gained the electrons that Mg lost, we say that oxygen was reduced. Therefore, Mg
did the reducing, making it the reducing agent. Mg would not have lost its electrons if oxygen
had not wanted them so much, therefore oxygen did the oxidizing making it the oxidizing
agent. The oxidizing agent becomes reduced when it removes electrons from the other reagent
in the redox reaction, while the reducing agent becomes oxidized as its electrons are lost.
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Each atom in a molecule is assigned an oxidation number, or the charge it would adopt if the
electrons were actually not shared in the bond but instead were transferred completely. Below
is a table of general guidelines to follow when determining the oxidation numbers of atoms in a
compound.
The oxidation number for an element in a binary ionic compound has a value based in reality
because it usually is equal to the ionic charge. Oxidation numbers for atoms in the polyatomic
ions or covalent compounds are not reality based because these atoms are, in fact, sharing their
electrons. But it does provide us with a guide because in a redox reaction, the oxidation
numbers of the species change, and it is most important to monitor these species and their
changes.
You should be able to identify the species that was oxidized and the species that was reduced
and label the oxidizing reagent and the reducing reagent.
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Reactions Types:
Decomposition – 1 chemical becomes 2 chemicals
AB+C
Combination – 2 chemicals become 1 chemical
A+BC
Single Replacement – 1 chemical replaces another chemical in a compound
A + BC  B + AC
Double Replacement – cations switch compounds
AB + CD  AD + CB
Combustion – a chemical is combined with oxygen in a rapid reaction.
A + O2  AxOy (+ Z…)
Redox – at least 2 chemicals in a reaction change their oxidation number/charge
Examples:
Decomposition
NH4NO2(s)  N2(g) + H2O(g)
Combination
P4(s) + O2(g)  P4O10(s)
Single Replacement
Zn(s) + CuCl2(aq)  Cu(s) + ZnCl2(aq)
Double Replacement
NaCl(aq) + AgNO3(aq)  AgCl(s) + NaNO3(aq)
Combustion
C3H8(g) + O2  CO2(g) + H2O(g)
Redox
Ca(s) + S(s)  CaS(s)
Note that all reactions above, except for the double replacement reactions, can fall under the redox category.
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