Transcript H 2 SO 4

Chapter 4
Reactions in Aqueous Solution
Solutions
A solution is a homogeneous mixture composed of a solute
and a solvent. The solute is dissolved in the solvent.
A solute is the substance present in smaller amount in a
solution. It can also be thought of as the dissolved
material.
A solvent is the substance present in larger amount in a
solution. It can also be thought of as the dissolving
material.
A solution can be gaseous (air), solid (alloy), or liquid
(seawater). In this chapter, we will only discuss aqueous
solutions, in which the solute initially is a liquid or a solid
and the solvent is water.
Molarity and Dilution
Molarity (M) is the number of moles of the
solute per liters of solution.
Molarity (M) = moles of solute = mol
Liters of Solution
L
When Diluting use:
M1V1 = M2V2
Hydration
Water is a very effective solvent for ionic compounds. Although
water is an electrically neutral molecule, it has a positive (H
atoms) and negative (O atoms) region, or “positive and negative
poles”. This is why water is called a polar solvent.
When an ionic compound, such as NaCl, dissolves in water, the
three dimensional network of ions in the solid is destroyed. The
Na+ and the Cl- ions are separated from one another and they
undergo hydration.
Hydration is the process in which an ion is surrounded by water
molecules arranged in a specific manner. Let’s see this in
action.
NaCl(s)
H2O
+
Na
(aq)
+
Cl (aq)
The term Dissociation means that the compound breaks up into cations
and anions like in the above equation for salt.
Solid NaCl, salt, is an ionic compound and breaks up into Na+ and Cl-,
cations and anions when dissolved in water. The Na+ ions are
attracted to the negative electrode and the Cl- anions are attracted to
the positive electrode. This movement sets up an electric current that
is equivalent to the flow of electrons along a metal wire.
Because NaCl conducts electricity, we say that NaCl is an electrolyte.
Pure Water contains very few ions and
therefore we call it a nonelectrolyte.
The above equation also shows that all of
the salt has dissociated into ions and there
is no undissociated NaCl left over in the
solution. This would be the same as
saying that salt is a very strong electrolyte.
Ionic vs. Molecular Compounds
Dissolving in Water
• As we have already talked about in earlier chapters, Molecular Compounds
dissolve in water but they break apart into molecules floating around in the
water and therefore they do not have a charge. This is why they are called
Molecular Compounds!
• Ionic Compounds break apart into ions when they are dissolved in water. The
ions have a charge in the solution and are good electrolytes! Remember all
bases are ionic except ammonia!
• Exception: Acids and the weak base ammonia (NH3) are considered molecular
compounds but they do break apart into ions in water:
NH3 + H2O D NH4+ + OHHCl + H2O D H3O+ + Cl• Also NOTE: Dissolving in water does not make something a strong electrolyte
(think of sugar = molecular and dissolves in water, it is not a strong electrolyte!)
All solutes that dissolve in water fit into one of 2 categories:
electrolytes and nonelectrolytes.
An electrolyte is a substance that, when dissolved in water, results in
a solution that can conduct electricity.
A nonelectrolyte does not conduct electricity when dissolved in
water.
Always USE g arrow
Always USE D arrow
Strong Electrolyte Weak Electrolyte
HCl
CH3COOH
HBr
HI
HF
HNO3
HNO2
HClO4
H2SO4*
NH3
All 1A: (LiOH, NaOH,
KOH, RbOH, CsOH)
H2O*
2A: Ba(OH)2 & Sr(OH)2
(all of the above loose
Ca(OH)2
+ ion when
an
H
Ionic Compounds
dissociated)
•H2SO4 has 2 ionizable
*Pure water is an
H+ ions, the second
extremely weak
form, HSO4- is a weak
electrolyte.
electrolyte
Nonelectrolyte
(NH2)2CO (urea)
CH3OH (methanol)
C2H5OH (ethanol)
C6H12O6 (glucose)
C12H22O11 (sucrose)
The strong/weak parts
of this chart should be
memorized because it
will help you to
memorize your
strong/weak acids and
bases.
Ionization of Acids and Bases
We use the term Ionization in order to describe the separation of
acids and bases into ions.
In order to determine whether or not you have a strong acid or base,
you see whether the acid or base dissociates completely in water.
If it does completely ionize in water, then it is considered to be a
strong acid or base. This would also make it a Strong
Electrolyte. To symbolize a strong electrolyte you use a single
arrow g. Such as: HCl(aq) g H+(aq) + Cl-(aq)
When an acid or base does not completely ionize in water, it is a
weak acid/base. These are also called Weak Electrolytes.
Acetic acid is a weak acid and we represent the ionization of acetic
acid with a double arrow D to show that it is a reversible
reaction, or that the reaction can occur in both directions.
CH3COOH(aq) D CH3COO-(aq) + H+(aq)
Solubility and Precipitates
Solubility is a term that is used in order to describe the
amount of solute that will dissolve in a solvent.
Most mixtures have a certain amount of solute that will
dissolve inside of a solvent.
Once the maximum amount of solute particles have been
added to a solvent, the solute particles will no longer
dissolve in the solvent and they will turn into the solid form
of the particle.
This change from liquid to solid is called a precipitate.
If something is considered to be insoluble in water, then it
means it has a solubility of less than .01 M in water at 25
degrees Celsius.
Solubility Rules for Common Ionic
Compounds in Water at 25 oC
Soluble Compounds
Compounds containing alkali metal ions (Li+, Na+,
K+, Rb+, Cs+) and the ammonium ion (NH4+)
Nitrates (NO3-), bicarbonates (HCO3-), and chlorates
(ClO3-)
Halides (Cl-, Br-, I-)
Sulfates (SO42-)
Insoluble Compounds
Carbonates (CO32-)
Phosphates (PO43-)
Chromates (CrO42-)
Sulfides
(S2-)
Hydroxides (OH-)
Exceptions
Memorize these Rules!!
Halides of Ag+, Hg22+, and Pb2+
Sulfates of Ag+, Ca2+, Sr2+, Ba2+,
Hg2+, and Pb2+
Exceptions
Compounds containing alkali
metal ions and the ammonium
ion
Compounds containing alkali
metal ions, ammonium ion and
the Ba2+ ion
Molecular Equations and Ionic Equations
Pb(NO3)2(aq) + 2NaI(aq) g PbI2(s) + 2NaNO3(aq)
The above equation is considered a molecular equation because
the formulas of the compounds are written as though all
species existed as molecules or whole units. A molecular
equation is useful because it identifies the reagents, if we
wanted to bring about this reaction in the lab.
However, a molecular equation does not accurately describe
what actually is happening at the microscopic level.
An ionic equation is used to represent what is occurring on the
microscopic level. The ionic equation shows dissolved
species as free ions. The above molecular equation would
have an ionic equation such as:
Pb2+(aq) + 2NO3- (aq) + 2Na+(aq) + 2I-(aq)g PbI2(s) + 2Na+(aq) + 2NO3- (aq)
Spectator Ions and Net Ionic Equations
An ionic equation includes spectator ions.
Spectator ions are ions that are not included involved in the
overall reaction.
Spectator ions appear on both sides of the equation and are
unchanged in the chemical reaction, therefore they can be
canceled. In the previous ionic equation:
Pb2+(aq) + 2NO3- (aq) + 2Na+(aq) + 2I-(aq) g
PbI2(s) + 2Na+(aq) + 2NO3- (aq)
You would cancel the 2NO3- and the 2Na+ spectator ions. The
net ionic equation, which shows only the species that
actually take part in the reaction, would be:
Pb2+(aq) + 2I-(aq) g
PbI2(s)
Note: If everything in a net ionic equation is a spectator ion
(and cancels on both sides) then you say: No Reaction
Occurred!!
Acids and Bases
Arrhenius defined an acid as producing a H+ ion
when dissolved in water. He also defined a base as
producing an OH- ion when dissolved in water.
This definition was limiting because it only applied
to aqueous solutions.
The chemists Johannes Bronsted and Thomas Lowry
proposed a more broad definition. Their definition does
not require that an acid be in an aqueous solution and
includes more than just protons and hydroxide ions for
acids and bases.
A Bronsted-Lowry acid is a proton (H+) donor.
A Bronsted-Lowry base is a proton (H+) acceptor.
Hydrochloric Acid
Ammonia
Water
Water
Hydronium Ion
Ammonium Ion
Chloride Ion
Hydroxide Ion
Monoprotic, Diprotic and Triprotic Acids
When each unit of acid only produces one hydrogen ion upon ionization
(or hydronium ion), that acid is said to be a monoprotic acid.
Examples are: hydrochloric (HCl), nitric (HNO3) and acetic acid
(CH3COOH). These are common acids you should memorize.
HNO3(aq) g H+ + NO3-(aq) (Strong Acid/Electrolyte)
When each unit of acid produces 2 hydrogen ions upon ionization (or
hydronium ions), that acid is said to be a diprotic acid. Example:
Sulfuric Acid (H2SO4). This is also a common acid that you should
memorize.
H2SO4(aq) g H+(aq) + HSO4-(aq) (Strong Acid/Electrolyte)
HSO4-(aq) D H+(aq) + SO42-(aq) (Weak Acid/Electrolyte)
When each unit of acid produces 3 hydrogen ions upon ionization (or
hydronium ions), that acid is said to be a triprotic acid. Example:
Phosphoric Acid (H3PO4). This is also a common acid that you
should memorize.
Let’s try and represent the triprotic acid of phosphoric acid on our own.
Acid-Base Neutralization
A neutralization reaction is a reaction between an acid and a base.
Generally, aqueous acid-base reactions produce water and a salt, which is
an ionic compound made up of a cation other that H+ and an anion other
that OH- or O2-:
acid + base g salt + water
HCl(aq) + NaOH(aq) g NaCl(aq) + H2O(l)
All salts are strong electrolytes. The substance we know as table salt, NaCl,
is a familiar example.
However, since both the acid and the base are strong electrolytes, they are
completely ionized in solution. The ionic equation is:
H+(aq) + Cl-(aq) + Na+(aq) + OH-(aq) g Na+(aq) + Cl-(aq) + H2O(l)
Therefore, the net ionic equation is:
H+(aq) + OH-(aq) g H2O(l)
Both Na+ and Cl- ions are spectator ions.
If we had started the reaction with equal molar amounts of the acid and the
base, at the end of the reaction we would have only a salt and no leftover
acid or base. This is a characteristic of acid-base neutralization reactions.
Acid-Base Titration
In titration, a solution of accurately known concentration, called a
standard solution, is added gradually to another solution of
unknown concentration, until the chemical reaction between the
two solutions is complete.
The point at which the acid has completely reacted with or has been
neutralized by the base is called the equivalence point.
The endpoint is the point at which the solution should change in
color due to the indicator, which changes color at or near the
equivalence point.
Many times, solid sodium hydroxide is not pure because of its
absorbance of water from the air. Therefore, when we make up
solutions of sodium hydroxide in the lab, we are unsure of it’s
exact molarity and we must standardize the solution.
Standardizing the solution is when we use the acid-base titration
technique in order to verify what the molarity of the unknown
solution is.
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Diprotic Titrations
How many milliliters of a 0.610 M NaOH solution are
needed to neutralize 20.0 mL of a 0.245 M H2SO4
solution? The equation for the reaction is:
2NaOH(aq) + H2SO4(aq) g Na2SO4(aq) + 2H2O(l)
Notice that we need twice the amount of sodium hydroxide
in order to neutralize the diprotic sulfuric acid. This is
because:
H2SO4(aq) g H+(aq) + HSO4-(aq)
In the above equation, one mole of OH- ions will neutralize one
mole of the H+ ions.
HSO4-(aq) D H+(aq) + SO42-(aq)
In the above equation, one mole of OH- ions will neutralize one
mole of the H+ ions.
Therefore, we need a total of 2 moles of NaOH in order to neutralize
one mole of H2SO4
Let’s Finish this problem on our own….
Gas Forming Reactions
Many reactions release a gaseous product. Although a wide variety of these gasforming reactions occur, some of the most important gases produced in reactions are the
following:
Acid/Base Rxns that form Gases: S2- (Sulfides) and CO32-(Carbonate), HCO3(Bicarbonate) ions react with acids to form gases:
CO2 Gas formation: A 2 Step Reaction:
1 step – Any CO32- or HCO3- ion reacting with an acid gives Carbonic Acid (H2CO3):
Example: HCl(aq) + NaHCO3(aq) g NaCl(aq) + H2CO3(aq)
2 step – Carbonic Acid is unstable and will decompose immediately into CO2 gas and
liquid H2O:
Example: H2CO3(aq)g CO2(g) + H2O(l)
H2S Gas (smells like rotten eggs) formation: forms when an acid (like HCl) reacts
with a Metal Sulfide (like Na2S):
Example: 2HCl(aq) + Na2S(aq) g H2S(g) + 2NaCl(aq)
H2 Gas forms in a single replacement (REDOX) reaction when an acid reacts
with a metal:
Example: Ca(s) + 2HCl(aq) g CaCl2(aq) + H2(g)
Other Rxns that form Gases:
O2 formed in many ways, one example: Electrolysis of water
NO2 formed when in air when lightning hits to supply energy:2NO(g) + O2(g) g 2NO2(g)
Rules for Assigning Oxidation States
1.
2.
3.
4.
5.
6.
7.
8.
The oxidation state of an atom in an element is 0. Atoms in their elemental
form are 0, example H2 is the elemental form of Hydrogen therefore in the
H2 molecule, each H = 0 (all diatomic atoms are the same) or in P4 each P =
0 or in S8 each S = 0.
In a neutral molecule, the sum of the oxidation numbers of all the atoms must
equal 0.
The oxidation state of a monatomic ion is the same as its charge.
In its compound, fluorine is always assigned an oxidation sate of -1.
Oxygen is usually assigned an oxidation of -2 in its covalent compounds,
such as CO, CO2, SO2. Exceptions to this rule includes peroxides
(compounds containing the O22- group), where each oxygen is assigned an
oxidation state of -1, as in hydrogen peroxide (H2O2), and OF2 in which
oxygen is assigned a +2 oxidation state.
In its covalent compounds with nonmetals, hydrogen is assigned an oxidation
state of +1. For example HCl. When hydrogen is bonded to a nonmetal in a
binary compound, it is assigned an oxidation state of -1. For example LiH.
For an ion, the sum of the oxidation states must equal the charge of the ion.
For example, the sum of the oxidation states must equal -2 in CO32-.
When Halogens combine with Oxygen, then halogens have a + charge.
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Oxidation-Reduction Reactions Redox
Oxidation-Reduction Reactions are considered
electron-transfer reactions.
Also known as Redox Reactions.
Remember: LEO goes GER
Loss of Electrons means Oxidized & Gain of Electrons
means Reduced
If you say an element is Oxidized, then it is called a
Reducing agent because it donates electrons to
another element.
If you say an element is Reduced, then it is called an
Oxidizing agent because it accepts electrons from
another element.
Let’s take a look at the formation of calcium oxide (CaO)
from calcium and oxygen:
2Ca(s) + O2(g) -> 2CaO(s)
You should recognize this is a REDOX reaction by checking
the individual oxidation charges of each atom/ion and
verifying that Loss of Electrons & Gain of Electrons is
occurring:
Loss of electrons: 2Ca -> 2Ca2+ + 4eCalcium is therefore being Oxidized
Therefore Calcium is a good Reducing Agent!
Gain of electrons: O2 + 4e- -> 2O2While Oxygen is being Reduced!
Therefore Oxygen is a good Oxidizing Agent!
Remember:
LEO goes GER
Single Replacement Reactions are
one type of REDOX Reactions!
• Single Replacement Reactions are one type of
REDOX reactions (We will cover other types
of REDOX reactions later – but you should be
able to recognize a REDOX reaction by
checking the oxidation charges of atoms/ions).
• You must use the ACTIVITY SERIES to
determine if a single replacement reaction will
actually occur. The Activity Series for Metals
is on the right. The Activity Series for
nonmetals is on the left.
• An element can only replace another element
that is less active than itself. Sort of like the
top boss can kick out anyone below him and
use their office. Examples:
3Mg + 2 FeCl3 g 2 Fe + 3 MgCl2
Cl2 + 2KI g I2 + 2KCl
Cl2 + KF g No Reaction!
MOST
ACTIVE
Li
Rb
K
Ba
Sr
Ca
Na
Mg
Al
Mn
F2
Cl2
Zn
Cr
Fe
Cd
Br2
I2
Co
Ni
Sn
Pb
H2
Sb
Bi
Cu
Hg
LEAST
ACTIVE
Ag
Pt
Au
Redox Reactions In Net Ionic Equation Form
Cu(s) + 2NO3-(aq) + 4H+ -> 2NO2(g) + Cu2+(aq) + 2H2O(l)
2Ag+(aq) + Cu(s) -> 2Ag(s) + Cu2+ (aq)
2NO(g) + O2(g) -> 2NO2(g)
2Fe(s) + 3Cl2(g) -> 2FeCl3(s)
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