Download Matter and Measurement

Chemical Reactions
Chemistry is about reactions with molecules colliding and
forming new molecules.
A number of reactions can be classified as "types”, allowing
some general predictions on outcomes of reactions.
Understanding the mechanism of reactions is important to
our understanding of processes such as acid rain, corrosion,
stain removers.
Dissolution Reactions
Precipitation Reactions
Acid-Base Reactions
Oxidation-Reduction Reactions
Dissolution Reactions
In dissolution reactions two or more compounds disperse
into each other to form a homogenous phase.
The starting compounds could be different phases (e.g. a
solid and a liquid), but the outcome of dissolution is a
homogenous phase; a SOLUTION
In dissolution reactions, the compound of lower
concentration is called the SOLUTE and the higher
concentration component is the SOLVENT.
During dissolution the solvent interacts with the solute such
that for the solute, the interactions between the solute and
solvent dominate over the solute-solute interactions and
solvent-solvent interactions.
Dissolution reactions are considered to be intermediate
between a chemical and physical process.
In terms of it being considered to be a chemical process,
solute-solute interactions are broken up and replaced by
solute-solvent interactions.
On the other hand the solution that results cannot be
expressed as a chemical formula and hence the outcome of
dissolution cannot be represented as a typical chemical
equation.
To write an equation for a dissolution reaction the solvent is
left out and the change in state of the solute denoted.
For example: dissolving sucrose in water
C12H22O11(s)  C12H22O11 (aq)
s - solid
aq - aqueous solution.
Note: for dissolution reactions, the solvent need not be water,
nor necessarily a liquid
Other examples of common liquid solvents are , but benzene
(C6H6), acetone (CH3COCH3), carbon tetrachloride (CCl4),
methanol (CH3OH)
Dissolution of Ionic Compounds
Most ionic compounds dissolve easily in water.
As we have seen ionic compounds, like NaCl, have rigid
lattices defined by the oppositely charged ions.
For the ionic compound to dissolve in water, the water
molecules must overcome the strong interactions that exist
between the oppositely charged species so that ion-water
(solute-solvent) interactions dominate over ion-ion
interactions (solute-solute)
The negative end of the water molecule (O) interacts with the
positive ions in the crystal and the positive end of water (H)
interacts with the negative end
When a water molecule encounters an ion it orients itself so
that the appropriate "side" of the water molecule interacts
with the ion (for negative ions, the H points toward the ion
and for positive ions, the O).
Having oriented itself in this way, the water molecule
essentially pulls this ion out of the crystal lattice.
Other water molecule surround this one ion and screens the
ion from the oppositely charged ions in the crystal.
Hence the water molecules, solvate the ion, and the solvated
ion then moves through the solution.
By this process the ionic compound dissolves in water and is
said to DISSOCIATE INTO ITS IONS.
Dissolution reactions for ionic compounds are written as:
NaCl(s)  Na+(aq) + Cl-(aq)
Dissolution of Covalent compounds
Neutral molecules do not have charges by which they can
interact with the solvent.
However, they can interact with the solvent through their
polarity.
Since the polarity of a molecule is due to a charge separation
within the molecule, solute molecules that are polar can
dissolve in polar solvents in much like the way ions dissolve
in water (a polar solvent).
For polar solutes, the polar solvent molecules orient
themselves around the solute molecules so that the more
positive end of the solvent is oriented towards the more
negative solute molecule and the more negative end of the
solvent orients itself towards the more positive end of the
molecule.
In this way the solute molecules are solvated replacing the
solute-solute interactions by solute-solvent interactions.
The solute molecules remain intact, but each solute
molecule is solvated by solvent molecules.
Dissolution of molecular compounds can be written as:
C12H22O11(s)  C12H22O11 (aq)
For a solute molecule to go “into solution”, the solvent
molecules must solvate the solute molecule so that the
solute-solvent interaction dominate over the solute-solute
interactions.
Hence, solute molecules dissolve in solutions of the same
polarity - “like dissolves like”.
So polar solutes dissolve in polar solvents, non-polar
solutes dissolve in non-polar solvents
Soaps or surfactants are “designed” so that the soap
molecule dissolves in water, yet can interact with a nonpolar oil molecule in a grease stain
a) Dissolution of an
ionic compound
b) Dissolution of a
covalent compound
Solubilities
Ethanol and water dissolve in each other and are said to be
miscible.
As more ethanol is added, at some point the ethanol
concentration is larger than the water concentration and the
solute and solvent switch.
Water and ethanol have infinite solubilties in the other.
Solubility is defined as the amount of a solute that can
dissolve in a fixed amount of solvent, at a given temperature
Solubility varies with temperature - generally higher the
temperature larger is the solubility.
As NaCl is added to water, a point is reached when the NaCl
does dissolve, but remains as a solid in the salt solution.
The point at which the NaCl stops dissolving in the salt
solution, defines the solubility of NaCl in water, at that
temperature.
The solution is said to be saturated with NaCl.
For a liquid solute dissolved in a liquid solvent, at the
saturation point, a new layer is formed, with the new layer
contains the solute with some of the original solvent
dissolved.
Electrolytes and Non-Electrolytes
battery
+
+
bulb
-
deionized, pure, water
battery
+
Na+
Cl-
+
bulb
-
deionized water + NaCl
NaCl(s) --> Na+(aq) + Cl-(aq)
When the battery is turned on the Na+ ions flow toward the
negative plate (anode) and the Cl- ions to the positive plate
(cathode).
The flow of ions constitutes a current. The circuit is now
complete, current flows through the circuit, and the bulb
turns on.
NaCl is called an electrolyte.
Electrolyte : a compound which when dissolved in a solvent
dissociates to form ions in solution.
Typically electrolytes are ionic compounds since they
dissolve in solution to form ions.
Example: K2SO4
Some covalent compounds (like acids and bases) can
dissociate in solution to form ions.
Electrolytes are characterized as being strong or weak.
The strength of an electrolyte depends on the degree to
which the compound dissociates in water to form ions.
Hence ionic compounds like NaCl and K2SO4 which
dissociate completely in water are strong electrolytes.
Weak electrolytes do not dissociate extensively in waterconsequently the conductance of a solution of a weak
electrolyte in low.
Non-electrolytes do not dissociate in solution to form ions
and hence their solutions do not conduct electricity.
Precipitation Reactions
2 KI(aq) + Pb(NO3)2(aq) --> PbI2 (s) + 2 KNO3 (aq)
The reaction between the KI and Pb(NO3)2 results in the
formation of PbI2 which has a very low solubility in water and
forms a solid precipitate.
2 KI(aq) + Pb(NO3)2(aq) --> PbI2 (s) + 2 KNO3 (aq)
For compounds insoluble in water, the attraction between the
oppositely charged ions in the solid crystal are too strong to
be overcome by solvent water molecules.
KNO3, being soluble in water, exists in solution as K+ and
NO3- ions.
The reaction
2 KI(aq) + Pb(NO3)2(aq) --> PbI2 (s) + 2 KNO3 (aq)
is also called a METATHESIS reaction.
In a metathesis reaction atoms or groups of atoms are
switched.
In this example K+ and Pb2+ switch anions.
NOTE: Metathesis reactions do not have to result in
precipitation.
Precipitation reactions result when
1) an amount exceeding the compound’s solubility in a
particular solvent is added to the solvent
2) Removal of solvent - example by evaporation.
3) Changing the solvent - since solubilities vary from solvent
to solvent, changing the solvent can result in precipitation
4) Changing the temperature - solubilities vary with
temperature. Cooling a solution of a compound can result in
the compound precipitating out of solution.
Predicting Precipitation Reactions
To determine if a product of a reaction between two ionic
compounds will result in a precipitate being formed, check
the solubilities of the compounds formed by the reaction.
Example: If aqueous solutions of AgNO3 and NaCl are mixed
will a precipitate form? If yes, identify the precipitate.
AgNO3 (aq) + NaCl (aq) --> AgCl (?) + NaNO3 (?)
Looking at the table
AgCl is insoluble, NaNO3 is soluble
AgNO3 (aq) + NaCl (aq) --> AgCl (s) + NaNO3 (aq)
Net Ionic Equations
AgNO3(aq) + NaCl(aq) --> AgCl(s) + NaNO3(aq)
If we re-write the above equation in terms of the species that
actually exist in solution, the equation is:
Ag+(aq) + NO3-(aq) + Na+(aq) + Cl-(aq)
--> AgCl(s) + Na+ + NO3-(aq)
This is the complete IONIC equation
In writing the complete ionic equation we see that Na+ and
NO3- exist in the same form on both sides of the equation,
whereas the Ag+ and Cl- have reacted to form solid AgCl.
Both Na+ and NO3- are called SPECTATOR IONS since they
are present in solution but do not participate directly in the
chemical reaction.
They “provide” the Ag+ and Cl- ions which reacted, but they
themselves do not directly participate in the chemical
reaction.
Hence, the complete ionic equation can be re-written to
show only those species which are directly involved in the
chemical reaction
Ag+(aq) + Cl-(aq) --> AgCl(s)
This is the NET IONIC equation.
Note: Net ionic equations, as with any chemical equation
must be balanced, in terms of mass and charge.
So for a reaction between Pb2+(aq) and I-(aq) to form PbI2(s)
must be written as:
Pb2+(aq) + 2 I-(aq) --> PbI2(s)
Example: Write the net ionic equation for the precipitation
reaction that occurs when aqueous solutions of calcium
chloride and sodium carbonate are mixed.
First write the chemical formulas of the reactants
aqueous Calcium chloride: CaCl2(aq)
aqueous sodium carbonate: Na2CO3(aq)
Next, determine what the products of the reaction will be and
which product is the precipitate.
The products of this reaction are: NaCl and CaCO3.
From the solubility table, determine if NaCl and CaCO3 are
soluble in water or not.
Write the net ionic equation - make sure it is balanced in
terms of mass and charge.
Hence, the equation for the reaction is:
CaCl2(aq) + Na2CO3(aq) --> NaCl(aq) + CaCO3 (s)
CaCl2(aq) + Na2CO3(aq) --> 2 NaCl(aq) + CaCO3 (s)
Ca2+(aq)+ 2Cl-(aq) + 2 Na+ + CO32- (aq) -->
2 Na+(aq) + 2Cl-(aq) + CaCO3 (s)
Canceling the spectator ions, gives the net ionic equation
Ca2+(aq) + CO32- (aq) --> CaCO3 (s)
Check that the ionic equation is balanced in terms of mass
and charge
Acid-Base Reactions
Acids and bases are probably one of the more commonly
encountered compounds
Acids are found in fruit (citric acid), vinegar (acetic acid), in
our stomachs (HCl), and in acid rain (H2SO4, HNO3).
Examples of bases - aqueous solutions of ammonia (used in
household cleanser); antacids like “Tums” and “Rolaids”
Acids and bases are also electrolytes they dissociate in water
to form ions.
Arrhenius Acids and Bases
Arrhenius acids are compounds which in aqueous solution
dissociate to form H+ ions.
(The H+ ion is also referred to as a proton since a H+ ion does
not have an electron and the charge is due to the single
proton in the nucleus)
Hence acids are H+ donors or proton DONORS.
Examples:
HCl(aq) --> H+(aq) + Cl-(aq)
HNO3(aq) --> H+(aq) + NO3-(aq)
Acids like HCl and HNO3 are called MONOPROTIC acids
since every molecule of HCl or HNO3 produces one H+
Acids like HCl and HNO3 completely dissociate in water
These acids are are STRONG ACIDS, and hence also strong
electrolytes
H2SO4(aq) --> H+(aq) + HSO4-(aq)
The HSO4- formed can dissociate further producing a H+.
However, HSO4- is not as strong an electrolyte as H2SO4
and not all the HSO4- ions dissociate.
HSO4- (aq)
H+(aq) + SO42-(aq)
The fact not all the HSO4- ions in solution dissociate is
denoted by the arrows pointing in both direction
Dissociation of HSO4- is “incomplete”
H2SO4 is a strong electrolyte since it completely dissociates
in water and hence is a strong acid.
HSO4- is a weak electrolyte since it does not completely
dissociate in water and hence a weak acid.
An aqueous solution of H2SO4 contains H+, HSO4- and SO42ions.
H2SO4 is called a DIPROTIC acid, since each molecule of
H2SO4 can produce up to 2 H+ ions.
Polyprotic acids - produce more that 2 H+ /molecule of acid
Arrhenius Bases
An Arrhenius base is a compound that when dissolved in
water dissociates to produce an OH- (hydroxide) ion.
NaOH(s) --> Na+(aq) + OH-(aq)
Compounds that do not contain OH- can still be bases as
long as when the dissolve in water the chemical reaction that
results produces OH- ions.
For example, if NH3 is dissolved in water, it can react with the
water to produce OHNH3(g) + H2O(l)
NH4+(aq) + OH-(aq)
In this example, note that the reaction is not complete as
indicated by the double arrows.
NH3 is a weak electrolyte and hence a weak base.
Also, note that NH3 received a H+ from water, and so in this
reaction the water acts as an acid giving up an H+ to NH3.
Water itself can dissociate as follows:
H2O(l)
H+(aq) + OH-(aq)
producing both H+ and OH- ions.
However the extent to which pure water dissociates is
very,very small, and a very small number of water molecules
dissociate.
In fact, pure water is not considered to be an electrolyte and
does not conduct electricity.
Compounds, like water, that can exhibit both acidic and basic
properties are called AMPHOTERIC.
Other examples of amphoteric compounds are amino acids
which exhibit both acidic and basic properties.
Reactions between acids and bases
When solutions of an acid and a base are mixed, a
NEUTRALIZATION reaction occurs.
The products of the neutralization reaction have neither
acidic nor basic properties.
Example:
HCl(aq) + NaOH(aq) --> H2O(l) + NaCl(aq)
In general:
acid + base --> salt + water
The complete ionic equation for the reaction between HCl and
NaOH is
H+(aq) + Cl-(aq) + Na+(aq) + OH-(aq) -->
H2O(l) + Na+(aq) + Cl-(aq)
Therefore, the net ionic equation is
H+(aq) + OH-(aq) --> H2O(l)
Problem
Milk of magnesia is essentially Mg(OH)2 and is insoluble in
water. Adding HCl to a suspension of milk of magnesia
dissolves it leaving behind a clear solution.
Write the overall equation and the net ionic equation.
Overall equation
Mg(OH)2(s) + 2 HCl(aq) --> MgCl2(aq) + 2 H2O(l)
Complete ionic equation
Mg(OH)2(s) + 2H+(aq) + 2Cl-(aq) -->
Mg2+(aq) + 2Cl- (aq) + 2H2O(l)
Net ionic equation
Mg(OH)2(s) + 2H+(aq) -> Mg2+(aq) + 2H2O(l)