Reactions - Chapter 4 Aqueous Reactions and Solution Chemistry by pptfiles

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									Aqueous Reactions and Solution Chemistry (4)

Solutions

Solution = Homogeneous mixture

A solution is a Homogeneous mixture (a mixture that has uniform properties
through out the material)

Solvent/Solutes

A solution is composed of a solute dissolved in a solvent. Normally, the
solute is the lesser quantity and the solvent is the greater quantity.
There is one solvent, there can be one or more solutes.

Aqueous Solutions

Are solutions with water as the solvent

Molarity
An expression of concentration

Molar Concentration, also Molarity, symbol (M), is defined as

Molarity = moles of solute/ liters of solution

M = n/V        or      n = MV

Molarity is often more useful since chemicals react on a mole stoichiometry basis and not
a weight basis.

Writing Unit Factors
The unit factor from molarity is moles/liter or it can also be written moles/1,000 mL

Solving Molarity Problems
The concept of molarity allows you to associate moles of solute with volume of solution.
Using molar mass also allows a connection of mass of solute with volume of solution.

How many grams of NaCl is required to make 1.000 L of 1 M solution?

1.000 L x 1 mole/L x 58.44 g/mole = 58.44 grams NaCl
149.1 grams of KCl was added to enough water to make 1.000 L of solution. What is the
Molarity?

(149.1 g KCl/74.55 g/mol)/ 1.000L = 2.000 moles/L = 2.000 M KCl

What volume of 8.00 M KCl solution is required to make 500 mL of 2.00 M KCl
solution?

(2.00 M x 0.500 L) = (8.00 M x ? L)

(2.00 M x 0.500 L)/8.00 M = 0.125 L = 125 mL

Concentration of Ions
When an ionic compound dissolves, the Molarity represents the concentration of the
compound in solution. It does not necessarily represent the concentration of ions in
solution. The concentration of ions will depend on the number of ions produced by the
compound.

Dilution
Since concentration (molarity) times volume gives moles [MV = n], this can be used to
calculate dilution volumes using

M1V1 = n = M2V2

Stoichiometry Involving Concentration
Titration is commonly used to determine the concentration of an unknown
compound. A solution of known concentration is reacted with a solution of
unknown concentration, with the amount added is controlled until the
reaction is just complete. This provides enough information to calculate the
concentration of the unknown.

M1V1 (n2/n1) = M2V2
or

(M1V1/n1) = (M2V2/n2)

The Nature of Aqueous Solutions: Strong and Weak Electrolytes

Dissolving
The process of mixing a solute into a solution.

Ionic compounds, acids, and bases dissociate or ionize in the process of
dissolving in water,
Covalent (molecular) compounds generally (except for acids and bases) do
not dissociate in water.

Dissociation / Ionization

The process where an ionic compound, acid, or base separates into or reacts
to form individual cations and anions.
Example (ionic compound): Table salt is a solid consisting of sodium and chlorine;
NaCl(s)
As table salt is dissolved in water, the solid separates into individual ions.
NaCl(s)  Na+(aq) + Cl-(aq)

Example (acid): Vinegar is a solution consisting of acetic acid in water. Acetic acid is a
liquid that readily mixes with water. A minority of the acetic acid molecules will ionize
in water producing the hydrogen ion and the acetate ion.
CH3COOH(aq)  CH3COO-(aq) + H+(aq)

Example (base): Ammonia, NH3, is a weak base that produces hydroxide ions by reaction
with water.
NH3(aq) + H2O(l)  NH4+(aq) + OH-(aq)

Electrolytes

are compounds that in solution will conduct electricity. This occurs because
the compound forms ions in solution.
Electrolytes are either ionic compounds, acids or bases.

Strong Electrolytes

carry a large current, this is when compounds form ions readily, i.e. they
completely ionize in water.
Strong Electrolytes include all ionic compounds and strong acids

Weak Electrolytes

carry little current, this is when compounds don't form ions readily, they
dissociate only partially in water.
Weak electrolytes include weak acids and weak bases.

Nonelectrolytes

do not carry electrical current. These are compounds that do not form ions.

Nonelectrolytes include covalent compounds that are not acids nor bases.

Solubility
Solubility is the maximum amount of solute that can dissolve in solution at
a given temperature; usually given in grams solute per 100 mL of solvent.

Soluble: when a compound has an appreciable solubility in water.
Insoluble: when a compound has a minor or negligible solubility in water.
Solubility is temperature dependent: solubility of solids and liquids tend to
increase with increasing temperature; solubility of gases decrease with
increasing temperature.

Types of Chemical Reactions
Precipitation Reactions
Acid-Base Reactions
Oxidation-Reduction Reactions

Precipitation Reactions
Precipitation reactions are double replacement reactions in which one of the products is
insoluble.

Double replacement reaction: Two compounds exchange anions, or cations. (two cations
exchange anions, or two anions exchange cations) [Reactants and products consist of two
ionic compounds with the ions exchanged with each other]

       AX + BY  AY + BX

       Example: NaI(aq) + AgNO3(aq)  AgI(s) + NaNO3(aq)

Precipitate
One product will be listed as (s), which means that it is solid (a precipitate).
Predicting Reaction Products
Solubility rules are used to predict if a reaction will occur.

The products are predictable since the ionic charge of the metals will not change.
The precipitation product is neutral.
Generally, the precipitate will contain one type of cation and one type of anion.

General form: AX + BZ  AZ + BX

NaBr(aq) + AgC2H3O2  NaC2H3O2(aq) + AgBr(s) [This reaction occurs since one
product is insoluble]

NaCl(aq) + NH4NO3  NR {NaNO3(aq) + NH4Cl(aq)} [This reaction does not occur
since both products are soluble]

2NaCl(aq) + Pb(C2H3O2)2(aq)  NaC2H3O2(aq) + PbCl2(s) [This reaction occurs because
one product is insoluble]

NaCl(aq) + KNO3(aq)  NR (NaNO3(aq) + KCl(aq)) [This reaction does not occur
because both products are soluble]

BaBr2(aq) + Ag2SO4(aq)  2AgBr(s) + BaSO4(s) [This reaction occurs because both
products are insoluble]

Solubility Rules
Compounds containing the following ions are generally soluble.

       Nitrate ion, NO3-
       Alkali metal ions and ammonium ion; Li+, Na+, K+, Rb+, Cs+, NH4+
       Acetate ion, C2H3O2-
       Most halide ions, Cl-, Br-, I-, [Halides of Ag+, Hg22+, and Pb2+ are insoluble]
       Sulfate ion, SO42-, [Sulfates of Sr, Ca, Ba, and Pb are marginally soluble]
Compounds containing the following ions are generally insoluble in water.
       Hydroxide ion, OH- [Hydroxides of alkali metal ions and ammonium ion are soluble, Ca, Sr, Ba
        are slightly soluble]
       Carbonate ion, CO32-          [alkali metal ions and ammonium ion are soluble]
       Phosphate ion, PO43-          [alkali metal ions and ammonium ion are soluble]
       Sulfide ion, S   2-
                              [alkali metal ions and ammonium ion are soluble]

Describing Reactions in Solution

Molecular Equation
The molecular equation shows the molecules or formula's of compounds that react and
are formed.

Na2CrO4(aq) + CaS(aq)  CaCrO4(s) + Na2S(aq)

Complete Ionic Equation
The Complete Ionic Equation shows the form of all species in solution. All aqueous ionic
compounds get separated into ions.

2Na+(aq) + CrO42-(aq) + Ca2+(aq) + S2-(aq)  CaCrO4(s) + 2Na+(aq) + S2-(aq)

Spectator Ions
Ions that do not participate in the reaction are spectator ions. These ions show up as both
reactants and products in the complete ionic equation.

Net Ionic Equation
The Net Ionic Equation shows only the species that directly participate in the reaction.
Write the complete ionic equation without the spectator ions to get the net ionic equation.

CrO42-(aq) + Ca2+(aq)  CaCrO4(s)

Acid-Base Reactions

Arrhenius Acids and Bases
Arrhenius Acid is a substance that ionizes in water to produce hydrogen
ions. These compounds consist of molecules with a hydrogen attached by a
polar bond.

Examples: HCl, HF, H2SO4, H3PO4, CH3COOH (HC2H3O2)
HCl(g) + H2O(l)  H+(aq) + Cl-(aq) + H2O(l)  H3O+(aq) + Cl-(aq)

In water, the hydrogen ion, H+, combines with water to form the hydronium
ion, H3O+. The use of either form is acceptable.

Arrhenius Base is a substance that ionizes in water to produce hydroxide
ions.

Examples: NaOH, Ca(OH)2, KOH, Mg(OH)2, NH3, amines
Polyprotic Acids
Acids can classified by the number of protons they can release in solution.
Acids with 2 or more acidic protons are called polyprotic acids.
Usually only the first dissociation step contributes to the concentration of
H+.

Monoprotic acids have one ionizable proton (HCl)

Diprotic acids have two ionizable protons (H2SO4)

Triprotic acids have three ionizable protons (H3PO4)

Binary Acids

Hydrogen with a nonmetal element (most but not all). Examples: HCl, HBr, HI, HF, H2S
Not Acids: H2O, NH3

Oxyacids

Many common acids are oxyacids, in which the acidic proton is attached to
an oxygen atom [contains hydrogen, oxygen, plus another
element{hydrogen with a polyatomic anion}]. Examples: H2SO4, H3PO4

Inorganic Acids

Inorganic acids are the binary acids and oxyacids.

Organic Acids (carboxylic acids)

Organic acids have a carbon backbone with hydrogen and a carboxyl group
(-COOH)

Acid Strength

Strong Acid

A strong acid is an acid that is almost all in the dissociated/ionized form.
The dissociation equilibrium lies far to the right.

A strong acid yields a weak conjugate base. (low affinity for a proton)
[much weaker base than water]

Weak Acid
A weak acid has a dissociation equilibrium that lies far to the left.
Most of the acid is in the undissociated/unionized form.

A weak acid yields a relatively strong conjugate base.

The conjugate base is stronger than water.

Strong Bases

Strong bases dissociate essentially completely (LiOH, NaOH, KOH, RbOH,
CsOH)

Weak Bases

dissociate or forms ions only partially in water

Neutralization Reactions
An Arrhenius acid neutralizes an Arrhenius base to form a salt (ionic
compound) and water.

HCl(aq) + NaOH(aq)  NaCl(aq) + H2O(l)
Acid + base         salt     + water



The net ionic equation of these neutralization reactions is generally:
H+(aq) + OH-(aq)  H2O(l)

Bronsted Acid-Bases

Bronsted-Lowry Acids and Bases
(A more general, inclusive, definition)
Bronsted Acid is a proton donor                      (proton is a hydrogen ion)

Bronsted Base is a proton acceptor [neutralizes the acid by removing its proton]
                                             [expands the definition of base]

Memory Device

BAAD: Base Accepts, Acid Donates

Conjugate acid
Conjugate acid is the species formed when the base accepts the proton.
Example: NH4+ is the conjugate acid when NH3 accepts a proton.

Conjugate base
The species that remains from the Bronsted acid species after the proton is
donated.
Example: F- is the conjugate base when HF donates a proton.

Conjugate acid-base pair
A Bronsted acid and its conjugate base. The acid-base pair differs by only
one proton.
Example: H2S and HS- are a conjugate acid-base pair.

A Bronsted Acid-Base neutralization reaction is the reaction of an acid and
base, this produces a weaker acid and base.
The hydroxide ion can be assumed to react completely with any acid (strong or weak).

Acid reactions with sulfides and Carbonates
These reactions produce gases; H2S and CO2 respectively.
Examples:
CaCO3(s) + 2HCl(aq)  CaCl2(aq) + CO2(g) + H2O(l)
CaCO3(s) + H2SO4(aq)  CaSO4(s) + CO2(g) + H2O(l)
Fe2S3(s) + 6HCl(aq)  2FeCl3(aq) + 3H2S(g)

Water as Acid and Base

Amphoteric (substance)

A substance is amphoteric if it can behave as either an acid or base.

Water is the most common amphoteric substance, as shown in the
autoionization of water where water is both acid and base in one reaction.
2H2O(l)  H3O+(aq) + OH-(aq)

The equilibrium expression is
Kw = [H3O+][OH-] = [1.0 x 10-7 M][1.0 x 10-7 M] = 1.0 x 10-14
Where the equilibrium constant, Kw, is called the ion-product constant or
dissociation constant of water
pH Scale

A way of expressing the concentration of the hydrogen ion is by using the
pH scale
Where
pH = -log[H+]

This also means that [H+] = 10-pH

Significant Figures: The number of decimal places in the log is equal to the
number of significant figures in the original number.

Similar log scales are used for other quantities including hydroxide.
pOH = -log[OH-]
Where pH + pOH = 14.

Calculating the pH of strong acid solutions

Strong acids are essentially completely dissociated. Use this information in
identifying the major species in solution. Identifying all the major species
will aid in determining the concentration in solution.

pH of Strong bases

[H+] is calculated using Kw (1.0 x 10-14) = [H+][OH-]

Electrochemistry: (redox reactions)
Electrochemistry is the study of the interchange of chemical and electrical
energy

Oxidation-Reduction Reactions
Oxidation-Reduction Reactions (or Redox Reactions) involve the transfer of electrons
between species.

Oxidation States
Oxidation States (or oxidation numbers) provide a way to keep track of electrons in
Oxidation-Reduction Reactions.

For covalent bonds with identical atoms, the electrons are shared between the atoms.
For two different atoms, the electrons are assigned to the atom that is more
electronegative (has a stronger attraction for electrons).

For ionic compounds, the oxidation state equals the ionic charge for monoatomic ions.

Assigning Oxidation States
      The oxidation state of an atom in an element is 0.

      The oxidation state of a monoatomic ion is the same as its charge.

      The sum of oxidation states for a neutral compound must be zero.

      The sum of oxidation states for a polyatomic ion equals the ionic charge.

      Group I metals, in compounds, have an oxidation state of +1.

      Group II metals, in compounds, have an oxidation state of +2.

      The following rules are hierarchical, the first listings has priority.

      The oxidation state of fluorine in compounds is always assigned a value of –1.

      The oxidation state of hydrogen in compounds with nonmetals is assigned a value
       of +1.

      The oxidation state of hydrogen in combination with metals is –1.

      The oxidation state of oxygen in compounds is usually assigned a value of –2,
       except for peroxides (O22-) where oxygen is assigned a value of –1; superoxide
       (O21-) where the oxygen is assigned a value of –1/2; or when combined with
       fluorine the oxygen value is variable.

      Group 7A nonmetals, Halogens (except fluorine) in compounds, generally are –1
       except when combined with oxygen or fluorine.

      Group 6A nonmetals, in compounds, has an oxidation state of –2.

      Group 5A nonmetals, in compounds, has an oxidation state of –3.

The Characteristics of Oxidation-Reduction (redox) Reactions

Oxidation-reduction (redox) reactions involves transfer of electrons from the
reducing agent to the oxidizing agent.

Oxidation and Reduction have to occur simultaneously.
Sometimes redox are obvious where the ionic charge of an element changes,
ions are formed or removed.

Sometimes, such as combustion reactions, it is less obvious because all
reactants and products are neutral covalent compounds.
Redox is a combined word from the two words:
      reduction and oxidation

Oxidation

Involves an increase in oxidation number/state (a lose of electrons)

Reduction

Involves a decrease in oxidation number/state (a gain of electrons)

Memory Devices

      OIL RIG: Oxidation Involves Loss; Reduction Involves Gain
      LEO says GER: Lose Electrons, Oxidized; Gain Electrons, Reduced

Oxidizing Agent (electron acceptor)

Oxidizing Agent (electron acceptor) promotes oxidation in another species.
The Oxidizing Agent is reduced itself.

Reducing Agent (electron donor)

Reducing Agent (electron donor) promotes reduction in another species. The
Reducing Agent is oxidized.

Half Reactions

Redox reactions can be broken into half reactions; one involving oxidation
and one involving reduction.
A Half Reaction shows either an oxidation reaction alone or a reduction
reaction alone.

The half reactions shows the electrons transferred; as a product for oxidation
reactions and as a reactant for reduction reactions.
Half-reactions are balanced for elements and charge.

Single Displacement Reactions (Oxidation of Metals by Acid or Salt)

Balancing Oxidation-Reduction Reactions
For any redox reaction to be balanced, the number of electrons gained and
the number of electrons lost must be equal.

Redox reactions can occur in an acidic environment, an alkaline (basic)
environment, or a neutral environment.

Balancing these types of reactions depends on the type of environment in
which the reaction occurs.

Half-reaction method in Aqueous Solutions
      Assign oxidation states to each element on each side of the reaction

      Separate the overall reaction into separate half reactions

      Balance each half-reaction

      Balance all elements except hydrogen and oxygen

      Balance oxygen using H2O

      Balance hydrogen using H+

      Balance the charge using electrons

      Multiply one or both balanced half-reactions by an integer to equalize the number
       of electrons transferred in both half-reactions

      Add the half-reactions and cancel any species which appears as both reactant and
       product.

      For basic solutions add OH- to both sides of the equation to neutralize any H+ into
       H2O, then remove any excess water so water only appears at most on one side of
       the equation.

      Check that all elements and charges are balanced

Balancing Example
Problem: Balance the following reaction in acidic solution:
Cr2O72-(aq) + NO2-(aq)  Cr3+(aq) + NO3-(aq)

Solution:

      Assign oxidation states to each element on each side of the reaction

Cr2O72-(aq) + NO2-(aq)  Cr3+(aq) + NO3-(aq)
+6 -2        +3 -2       +3        +5 -2

      Separate the overall reaction into separate half reactions

Oxidation: NO2-(aq)  NO3-(aq)

Reduction: Cr2O72-(aq)  Cr3+(aq)

      Balance each half-reaction

      Balance all elements except hydrogen and oxygen

Oxidation: NO2-(aq)  NO3-(aq)

Reduction: Cr2O72-(aq)  2Cr3+(aq)

      Balance oxygen using H2O

Oxidation: NO2-(aq) + H2O  NO3-(aq)

Reduction: Cr2O72-(aq)  2Cr3+(aq) + 7H2O

      Balance hydrogen using H+

Oxidation: NO2-(aq) + H2O(l)  NO3-(aq) + 2H+(aq)

Reduction: Cr2O72-(aq) + 14H+(aq)  2Cr3+(aq) + 7H2O(l)

      Balance the charge using electrons

Oxidation: NO2-(aq) + H2O(l)  NO3-(aq) + 2H+(aq) + 2e-

Reduction: Cr2O72-(aq) + 14H+(aq) + 6e-  2Cr3+(aq) + 7H2O(l)

      Multiply one or both balanced half-reactions by an integer to equalize the number
       of electrons transferred in both half-reactions

Oxidation (x3): 3NO2-(aq) + 3H2O(l)  3NO3-(aq) + 6H+(aq) + 6e-

Reduction: Cr2O72-(aq) + 14H+(aq) + 6e-  2Cr3+(aq) + 7H2O(l)
      Add the half-reactions and cancel any species, which appears as both reactant and
       product.

Add Half-reactions
Cr2O72-(aq) + 3NO2-(aq) + 3H2O(l) + 14H+(aq) + 6e-  2Cr3+(aq) + 3NO3-(aq) + 7H2O(l)
+ 6H+(aq) + 6e-
Cancel duplicate species
Cr2O72-(aq) + 3NO2-(aq) + 8H+(aq)  2Cr3+(aq) + 3NO3-(aq) + 4H2O(l)

      For basic solutions add OH- to both sides of the equation to neutralize any H+ into
       H2O, then remove any excess water so water only appears at most on one side of
       the equation.

      Check that all elements and charges are balanced

Cr2O72-(aq) + 3NO2-(aq) + 8H+(aq)  2Cr3+(aq) + 3NO3-(aq) + 4H2O(l)

Predicting Spontaneous Redox Reactions
In single replacement reactions, one metal replaces another metal from a
compound. The metal that has a stronger tendency to become the ionic
compound is more active than the metal that becomes metallic. Activity is
the tendency of a metal to loose an electron to become ionic.

A combination of an oxidation reaction higher on the list combined with a
reduction reaction from lower on the list is spontaneous.

Each half reaction is associated with an electrical potential (i.e. voltage). A
combination of reactions that creates a positive voltage is spontaneous.
When a reaction is written in the reverse direction, it changes the sign of the
voltage.

The Activity Series [Don't Memorize]
Half-Reactions (By decreasing Oxidation ability)
       Most Easily Oxidized [Most Reactive]
       Li(s)  Li+(aq) + e-                                                   ξo = 3.05 V
       K(s)  K+(aq) + e-                                                     ξo = 2.93 V
       Ca(s)  Ca2+(aq) + 2e-                                                 ξo = 2.87 V
       Na(s)  Na+(aq) + e-                                                   ξo = 2.71 V
       Mg(s)  Mg2+(aq) + 2e-                                                 ξo = 2.37 V
       Al(s)  Al3+(aq) + 3e-                                                 ξo = 1.66 V
       Mn(s)  Mn2+(aq) + 2e-                                                 ξo = 1.18 V
       Zn(s)  Zn2+(aq) + 2e-                                                 ξo = 0.76 V
       Cr(s)  Cr3+(aq) + 3e-                                                 ξo = 0.74 V
       Fe(s)  Fe2+(aq) + 2e-                                                 ξo = 0.44 V
       Ni(s)  Ni2+(aq) + 2e-                                                 ξo = 0.23 V
       Sn(s)  Sn2+(aq) + 2e-                                                 ξo = 0.14 V
       Pb(s)  Pb2+(aq) + 2e-                                                 ξo = 0.13 V
       H2(g)  2H+(aq) + 2e-                                                   ξo = 0.0 V
       Cu(s)  Cu2+(aq) + 2e-                                                ξo = -0.34 V
       Ag(s)  Ag+(aq) + e-                                                  ξo = -0.80 V
       Hg(l)  Hg2+(aq) + 2e-                                               ξo = -0.854 V
       Pt(s)  Pt2+(aq) + 2e-                                                      ξo = -?
       Au(s)  Au3+(aq) + 3e-                                                ξo = -1.50 V
Least Easily Oxidized [Least Reactive]

Will a Metal Dissolve in Acid?
If a metal is higher on the activity series list than Hydrogen, then it will
dissolve in acid.
If the electric potential for the oxidation of the metal is positive, then it will
dissolve in acid.

Acid-Base Titrations
A titration (volumetric analysis) involves precisely adding a measured volume of a
known concentration of a reactant (titrant) to a measured volume of an unknown
concentration of another reactant (analyte).

This is done until a chosen endpoint is reached. The endpoint is often marked by the
change of color of an indicator chemical.

The indicator is chosen to be close to the stoichiometric point or equivalence point.

For a good titration:
The reaction occurring must be known and relatively fast
The stoichiometric point must be marked
The volumes must be measured accurately.
Phenolphthalein is usually the chosen indicator for an acid-base titration.

The titrant is usually a strong acid or base.

Titration Calculations
Titration calculations are conducted identically to stoichiometry
calculations.

Reactions conducted in solution use solution concentration in calculations of
how much of one reactant will react with another reactant.

MV yields moles of solute.

For a reaction of the form
n1Reactant1 + n2Reactant2  Products
These relationships between reactant concentrations can be calculated using
M1V1/n1 = M2V2/n2

Rearranging slightly gives
M1V1(n2/n1) = M2V2
or
moles[1](reaction stoichiometry) = moles[2]
ations can be calculated using
M1V1/n1 = M2V2/n2
Rearranging slightly gives
M1V1(n2/n1) = M2V2
or
moles[1](reaction stoichiometry) = moles[2]

								
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