Aqueous Reactions and Solution Chemistry (4)

Chapter 4
[Brown/LeMay]
[Silberberg]

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) à 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) à CH₃COO^-(aq) + H⁺(aq)

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

Example (water): Water self-ionizes to a small degree (nonelectrolyte)
2H₂O(l) ßà H₃O⁺(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.

Equilibrium Condition

Characteristics - Forward/Reverse

Many reactions are reversible, such that the products can react together to form the reactants.

Equilibrium is dynamic with forward and reverse reactions continuing to occur but are balanced.

Dynamic Equilibrium: The rate of the forward reaction is equal to the rate of the reverse reaction.

Types of Chemical Reactions

Precipitation Reactions

Acid-Base Reactions

Oxidation-Reduction Reactions

Precipitation Reactions

Precipitation reactions are double replacement reactions [Exchange 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 à AY + BX

Example: NaI(aq) + AgNO₃(aq) à 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 à AZ + BX

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

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

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

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

BaBr₂(aq) + Ag₂SO₄(aq) à 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, SO₄^2-, [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, CO₃^2- [alkali metal ions and ammonium ion are soluble]

· Phosphate ion, PO₄^3- [alkali metal ions and ammonium ion are soluble]

· Sulfide ion, S^2- [Sulfides of alkali metal ions, ammonium ion, Ca, Sr, and Ba are soluble] [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) à 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) + CrO₄^2-(aq) + Ca²⁺(aq) + S^2-(aq) à CaCrO₄(s) + 2Na⁺(aq) + S^2-(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.

CrO₄^2-(aq) + Ca²⁺(aq) à CaCrO₄(s)

Acid-Base Reactions

Ch. 16 Sections 1-5 [Brown/LeMay]

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) à H⁺(aq) + Cl^-(aq) + H₂O(l) à 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) à NaCl(aq) + H₂O(l)
Acid + base à salt + water

The net ionic equation of these neutralization reactions is generally:
H⁺(aq) + OH^-(aq) à 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) à CaCl₂(aq) + CO₂(g) + H₂O(l)
CaCO₃(s) + H₂SO₄(aq) à CaSO₄(s) + CO₂(g) + H₂O(l)
Fe₂S₃(s) + 6HCl(aq) à 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) ßà H₃O⁺(aq) + OH^-(aq)

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

MCF08

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 K_w (1.0 x 10^-14) = [H⁺][OH^-]

Electrochemistry: (redox reactions)

Ch. 20 Sections 1-2 [Brown/LeMay]

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 (O₂^2-) where oxygen is assigned a value of –1; superoxide (O₂^1-) 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:
Cr₂O₇^2-(aq) + NO₂^-(aq) à Cr³⁺(aq) + NO₃^-(aq)

Solution:

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

Cr₂O₇^2-(aq) + NO₂^-(aq) à Cr³⁺(aq) + NO₃^-(aq)
+6 -2 +3 -2 +3 +5 -2

· Separate the overall reaction into separate half reactions

Oxidation: NO₂^-(aq) à NO₃^-(aq)

Reduction: Cr₂O₇^2-(aq) à Cr³⁺(aq)

· Balance each half-reaction

· Balance all elements except hydrogen and oxygen

Oxidation: NO₂^-(aq) à NO₃^-(aq)

Reduction: Cr₂O₇^2-(aq) à 2Cr³⁺(aq)

· Balance oxygen using H₂O

Oxidation: NO₂^-(aq) + H₂O à NO₃^-(aq)

Reduction: Cr₂O₇^2-(aq) à 2Cr³⁺(aq) + 7H₂O

· Balance hydrogen using H⁺

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

Reduction: Cr₂O₇^2-(aq) + 14H⁺(aq) à 2Cr³⁺(aq) + 7H₂O(l)

· Balance the charge using electrons

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

Reduction: Cr₂O₇^2-(aq) + 14H⁺(aq) + 6e^- à 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) à 3NO₃^-(aq) + 6H⁺(aq) + 6e^-

Reduction: Cr₂O₇^2-(aq) + 14H⁺(aq) + 6e^- à 2Cr³⁺(aq) + 7H₂O(l)

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

Add Half-reactions
Cr₂O₇^2-(aq) + 3NO₂^-(aq) + 3H₂O(l) + 14H⁺(aq) + 6e^- à 2Cr³⁺(aq) + 3NO₃^-(aq) + 7H₂O(l) + 6H⁺(aq) + 6e^-
Cancel duplicate species
Cr₂O₇^2-(aq) + 3NO₂^-(aq) + 8H⁺(aq) à 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

Cr₂O₇^2-(aq) + 3NO₂^-(aq) + 8H⁺(aq) à 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) à Li⁺(aq) + e^- ξ^o = 3.05 V

K(s) à K⁺(aq) + e^- ξ^o = 2.93 V

Ca(s) à Ca²⁺(aq) + 2e^- ξ^o = 2.87 V

Na(s) à Na⁺(aq) + e^- ξ^o = 2.71 V

Mg(s) à Mg²⁺(aq) + 2e^- ξ^o = 2.37 V

Al(s) à Al³⁺(aq) + 3e^- ξ^o = 1.66 V

Mn(s) à Mn²⁺(aq) + 2e^- ξ^o = 1.18 V

Zn(s) à Zn²⁺(aq) + 2e^- ξ^o = 0.76 V

Cr(s) à Cr³⁺(aq) + 3e^- ξ^o = 0.74 V

Fe(s) à Fe²⁺(aq) + 2e^- ξ^o = 0.44 V

Ni(s) à Ni²⁺(aq) + 2e^- ξ^o = 0.23 V

Sn(s) à Sn²⁺(aq) + 2e^- ξ^o = 0.14 V

Pb(s) à Pb²⁺(aq) + 2e^- ξ^o = 0.13 V

H₂(g) à 2H⁺(aq) + 2e^- ξ^o = 0.0 V

Cu(s) à Cu²⁺(aq) + 2e^- ξ^o = -0.34 V

Ag(s) à Ag⁺(aq) + e^- ξ^o = -0.80 V

Hg(l) à Hg²⁺(aq) + 2e^- ξ^o = -0.854 V

Pt(s) à Pt²⁺(aq) + 2e^- ξ^o = -?

Au(s) à Au³⁺(aq) + 3e^- ξ^o = -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 à 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]