Solutions and Aqueous Reactions Notes

Solution Concentration

• Homogeneous mixtures are called solutions.
• The component present in the greatest amount is the solvent.
• The component present in a lesser amount is the solute.
• A solution in which water is the solvent is an aqueous solution.
• Dilute solutions have a small amount of solute compared to solvents.
• Concentrated solutions have a large amount of solute compared to solvents.
• The greater the concentration, the greater the ratio of solute to solvent.

Solution Concentration: Molarity

• Molarity (M) is a common way to express solution concentration.
• Molarity is the amount of solute (in moles) divided by the volume of solution (in liters).

$Molarity (M) = \frac{amount \ of \ solute \ (in \ mol)}{volume \ of \ solution \ (in \ L)}$

Using Molarity in Calculations

• The molarity of a solution can be used as a conversion factor between moles of the solute and liters of the solution.
• For example, a 0.500 M NaCl solution contains 0.500 mol NaCl for every liter of solution.

Solution Dilution

• Solutions are often stored as concentrated stock solutions.
• To make solutions of lower concentrations from these stock solutions, more solvent is added.
• The amount of solute does not change, just the volume of solution.

$M<em>1V</em>1 = M<em>2V</em>2$

where

• $M_1$ = initial molarity
• $V_1$ = initial volume
• $M_2$ = final molarity
• $V_2$ = final volume

Solution Stoichiometry

• In aqueous reactions, we can use the volume and concentration of a reactant or product to calculate its amount in moles.
• We can then use the stoichiometric coefficients in the chemical equation to convert to the amount of another reactant or product in moles.
• The general conceptual plan for these kinds of calculations begins with the volume of a reactant or product.

What Happens When a Solute Dissolves?

• There are attractive forces between the solute particles holding them together.
• There are also attractive forces between the solvent molecules.
• When we mix the solute with the solvent, there are attractive forces between the solute particles and the solvent molecules.
• If the attractions between solute and solvent are strong enough, the solute will dissolve.

Charge Distribution in a Water Molecule

• There is an uneven distribution of electrons within the water molecule.
• The oxygen side of the molecule has a partial negative charge $(\delta^-)$, and the hydrogen side has a partial positive charge $(\delta^+)$.

Solute and Solvent Interactions in a Sodium Chloride Solution

• When sodium chloride is put into water, the attraction of $Na^+$ and $Cl^-$ ions to water molecules competes with the attraction among the oppositely charged ions themselves.

Dissolution of Ionic Compounds

• Each ion is attracted to the surrounding water molecules and pulled off and away from the crystal.
• When it enters the solution, the ion is surrounded by water molecules, insulating it from other ions.
• The result is a solution with free-moving, charged particles able to conduct electricity.

Electrolyte and Nonelectrolyte Solutions

• Substances that dissolve in water to form a solution containing ions will conduct electricity and are called electrolytes.
• Substances that dissolve in water to form a solution with no ions will not conduct electricity and are called nonelectrolytes.
• A solution of salt (an electrolyte) conducts electrical current. A solution of sugar (a nonelectrolyte) does not.
• Ionic substances, such as sodium chloride, that completely dissociate into ions when they dissolve in water, are strong electrolytes.
• Except for acids, most molecular compounds, for example sugar, dissolve in water as intact molecules and are nonelectrolytes.
• Acids ionize to varying degrees in water. Those that completely ionize are strong acids. Those that don’t are weak acids.

Acids

• Acids are molecular compounds that form $H^+$ ions when dissolved in water.
• Acids that completely ionize in water are called strong acids.
• Acids that do not completely ionize in water are called weak acids.

$HCl(aq) \rightarrow H^+(aq) + Cl^-(aq)$
$HF(aq) \rightarrow H^+(aq) + F^-(aq)$

Strong and Weak Electrolytes

• Strong electrolytes are substances that dissolve completely as ions.
  • Ionic compounds and strong acids.
  • Solutions are good conductors of electricity.
• Weak electrolytes are substances that dissolve mostly as molecules but partially as ions.
  • Weak acids.
  • Solutions conduct electricity, but not well.

The Solubility of Ionic Compounds

• When an ionic compound dissolves in water, the resulting solution contains not the intact ionic compound itself, but its component ions dissolved in water.
• However, not all ionic compounds dissolve in water.
• In general, a compound is termed soluble if it dissolves in water and insoluble if it does not.

Solubility of Salts

• If we mix solid $AgNO_3$ with water, it dissolves and forms a strong electrolyte solution.
• Silver chloride, on the other hand, is almost completely insoluble.
  • If we mix solid AgCl with water, virtually all of it remains as a solid within the liquid water.

When Will a Salt Dissolve?

• Whether a particular compound is soluble or insoluble depends on several factors.
• Predicting whether a compound will dissolve in water is not easy.
• The best way to do it is to conduct experiments to test whether a compound will dissolve in water and then develop some rules based on those experimental results.
  • We call this method the empirical method.

Solubility Rules

• Compounds Containing the Following Ions Are Generally Soluble
  • $Li^+, Na^+, K^+, and\ NH_4^+$: None
  • $NO<em>3^-, C</em>2H<em>3O</em>2^−$: None
  • $Cl^−, Br^−, I^−$: When these ions pair with $Ag^+, Hg_2^{2+}, or Pb^{2+}$ the resulting compounds are insoluble.
  • $SO<em>4^{2−}$: When $SO</em>4^{2−}$ pairs with $Sr^{2+}, Ba^{2+}, Pb^{2+}, Ag^+, or\ Ca^{2+}$, the resulting compound is insoluble.
• Compounds Containing the Following Ions Are Generally Insoluble Exceptions
  • $OH^−, S^{2−}$: When these ions pair with $Li^+, Na^+, K^+, or\ NH_4^+$ the resulting compounds are soluble. When $S^{2−}$ pairs with $Ca^{2+}, Sr^{2+}, or Ba^{2+}$, the resulting compound is soluble. When $OH^−$ pairs with $Ca^{2+}, Sr^{2+}, or Ba^{2+}$, the resulting compound is slightly soluble.
  • $CO<em>3^{2−}, PO</em>4^{3−}$: When these ions pair with $Li^+, Na^+, K^+, or\ NH_4^+$, the resulting compounds are soluble.

Precipitation Reactions

• Precipitation reactions are ones in which a solid forms when we mix two solutions.
  • Reactions between aqueous solutions of ionic compounds produce an ionic compound that is insoluble in water.
  • The insoluble product is called a precipitate.

Predicting Precipitation Reactions

1. Determine what ions constitute each aqueous reactant.
2. Determine formulas of possible products.
   • Exchange ions.
   • (+) ion from one reactant with (-) ion from other
   • Balance charges of combined ions to get the formula of each product.
3. Determine solubility of each product in water.
   • Use the solubility rules.
   • If the product is insoluble or slightly soluble, it will precipitate.
4. If neither product will precipitate, write no reaction after the arrow.
5. If any of the possible products are insoluble, write their formulas as the products of the reaction using (s) after the formula to indicate solid. Write any soluble products with (aq) after the formula to indicate aqueous.
6. Balance the equation.
   • Remember to change only coefficients, not subscripts.

Representing Aqueous Reactions

• A molecular equation is an equation showing the complete neutral formulas for each compound in the aqueous reaction as if they existed as molecules.
• In actual solutions of soluble ionic compounds, dissolved substances are present as ions. Equations that describe the nature of the dissolved species in solution are called complete ionic equations.

$2KOH(aq) + Mg(NO<em>3)</em>2(aq) → 2KNO<em>3(aq) + Mg(OH)</em>2(s)$

Writing a Complete Ionic Equation

• Aqueous strong electrolytes (soluble salts, strong acids, strong bases) are written as ions.
• Insoluble substances, weak electrolytes, and nonelectrolytes are written in molecule form.
  • Solids, liquids, and gases are not dissolved, hence written in molecule form.

$2K^+(aq) + 2OH^-(aq) + Mg^{2+}(aq) + 2NO<em>3^-(aq) → 2K^+(aq) + 2NO</em>3^-(aq) + Mg(OH)_2(s)$

Ionic Equation

• Some of the ions in solution appear unchanged on both sides of the equation in the complete ionic equation.
• These ions are called spectator ions because they do not participate in the reaction (soluble salts, strong acids, and strong bases).

Net Ionic Equation

• An ionic equation in which the spectator ions are removed is called a net ionic equation.

$Pb^{2+}(aq) + 2NO<em>3^-(aq) + 2K^+(aq) + 2Cl^-(aq) → PbCl</em>2(s) + 2K^+(aq) + 2NO_3^-(aq)$

The net ionic equation is

$Pb^{2+}(aq) + 2Cl^-(aq) → PbCl_2(s)$

Summarizing Aqueous Equations

• A molecular equation is a chemical equation showing the complete, neutral formulas for every compound in a reaction.
• A complete ionic equation is a chemical equation showing all of the species as they are actually present in solution: strong electrolytes are therefore represented as their component ions.
• A net ionic equation is an equation showing only the species that actually change during the reaction.

Acid–Base and Gas-Evolution Reactions

• Two other important classes of reactions that occur in aqueous solution are
  1. acid–base reactions, and
  2. gas-evolution reactions.
• Acid–base reaction:
  • Also called a neutralization reaction
  • An acid reacts with a base, and the two neutralize each other, producing water (or in some cases a weak electrolyte).
• In a gas-evolution reaction, a gas is produced, resulting in bubbling.
• In both acid–base and gas-evolution reactions, as in precipitation reactions, the reactions occur when the anion from one reactant combines with the cation of the other.
• Many gas-evolution reactions are also acid–base reactions.

Acid–Base Reactions

Arrhenius Definitions:

• Acid: Substance that produces $H^+$ in aqueous solution
  • In solution, $HCl$ bonds with water to produce the hydronium ion, $H_3O^+$.
  • Polyprotic acids contain more than one ionizable proton and release them sequentially.
  • The first ionizable proton is strong while subsequent ionizable protons are weak.
• Base: Substance that produces $OH^−$ ions in aqueous solution.

$NaOH \rightarrow Na^+ + OH^-(aq)$

• These reactions are called neutralization reactions because the acid and base neutralize each other’s properties.
• The net ionic equation for an acid–base reaction is

$H^+(aq) + OH^-(aq) \rightarrow H_2O(l)$

(As long as the salt that forms is soluble in water).

Acids and Bases in Solution

• Acids ionize in water to form $H^+$ ions.
  • More precisely, the $H^+$ from the acid molecule is donated to a water molecule to form hydronium ion, $H_3O^+$.
  • Most chemists use $H^+$ and $H_3O^+$ interchangeably.
• Bases dissociate in water to form $OH^-$ ions.
  • Bases, such as $NH_3$, that do not contain $OH^−$ ions, produce $OH^-$ by pulling off water molecules.
• In the reaction of an acid with a base, the $H^+$ from the acid combines with the $OH^−$ from the base to make water.
• The cation from the base combines with the anion from the acid to make the salt.

Acid–Base Titrations

• A titration is a laboratory procedure where a substance in a solution of known concentration (titrant) is reacted with another substance in a solution of unknown concentration (analyte).
• The equivalence point is the point in the titration when the $H^+$ and $OH^−$ from reactants are in their stoichiometric ratio and are completely reacted.
• An indicator is a dye whose color depends on the acidity or basicity of solution.

Gas-Evolving Reactions

• Some reactions form a gas directly from the ion exchange.
• Other reactions form a gas by the subsequent decomposition of one of the ion exchange products into a gas and water.

$H<em>2SO</em>4(aq) + Li<em>2S(aq) \rightarrow Li</em>2SO<em>4(aq) + H</em>2S(g)$
$HCl(aq) + NaHCO<em>3(aq) \rightarrow NaCl(aq) + H</em>2CO<em>3(aq)$$H</em>2CO<em>3(aq) \rightarrow H</em>2O(l) + CO_2(g)$

Types of Compounds That Undergo Gas-Evolution Reactions

Oxidation–Reduction Reactions

• The reactions in which electrons are transferred from one reactant to the other are called oxidation–reduction reactions, or redox reactions.
• Many redox reactions involve the reaction of a substance with oxygen.

$4Fe(s) + 3O<em>2(g) \rightarrow 2Fe</em>2O<em>3(s)$$2C</em>8H<em>{18}(l) + 25O</em>2(g) \rightarrow 16CO<em>2(g) + 18H</em>2O(g)$
$2H<em>2(g) + O</em>2(g) \rightarrow 2H_2O(g)$

Oxidation and Reduction

• Oxidation is the loss of electrons.
• Reduction is the gain of electrons.

$2Na(s) + Cl_2(g) \rightarrow 2NaCl(s)$

$Na(s) \rightarrow Na^+(g) + 1e^-(oxidation)$
$Cl_2(g) + 2e^- \rightarrow 2Cl^-(reduction)$

Redox Reaction

• Electron transfer does not need to be a complete transfer for the reaction to qualify as oxidation–reduction.
• Example:
• There is uneven sharing of electrons when hydrogen bonds to chlorine, resulting in an increase of electron density (reduction) for chlorine and a decrease in electron density (oxidation) for hydrogen.

Reactions of Metals with Nonmetals

• The reaction involves a metal reacting with a nonmetal.
• In addition, the reaction involves the conversion of free elements into ions.

$2Na(s) + Cl_2(g) \rightarrow 2NaCl(s)$

Oxidation States

• We need a method for determining how the electrons are transferred in reactions that don’t involve metal-nonmetal or combustion.
• Chemists assign a number to each element in a reaction called an oxidation state or oxidation number allowing them to determine the electron flow in the reaction, like electron bookkeeping.
• The oxidation number of an atom in a compound is the “charge” it would have if all shared electrons were assigned to the atom with the greatest attraction for those electrons.
• Oxidation states are imaginary charges assigned based on a set of rules.
  • They are written with charge before magnitude, for example +1 or -1.
• Ion charges are real, measurable charges.
  • They are written with magnitude before charge, for example $1^+$ or $1^-$.

Rules for Assigning Oxidation States

The following rules are in order of priority:

1. Free elements have an oxidation state = 0.
2. Monatomic ions have an oxidation state equal to their charge.
3. (a) The sum of the oxidation states of all the atoms in a compound is 0.
   (b) The sum of the oxidation states of all the atoms in a polyatomic ion equals the charge on the ion.
4. (a) Group 1A metals have an oxidation state of +1 in all of their compounds.
   (b) Group 2A metals have an oxidation state of +2 in all of their compounds.
5. In their compounds, nonmetals have oxidation states according to the table below. Nonmetals higher on the table take priority.

• When assigning oxidation states, keep these points in mind:
  • The oxidation state of any given element generally depends on what other elements are present in the compound (except groups 1A and 2A metals)
  • When following the hierarchy shown in rule 5, give priority to the element(s) highest on the list and then assign the oxidation state of the element lowest on the list using rule 3.
  • When assigning oxidation states to elements that are not covered by rules 4 and 5 (such as carbon), use rule 3 to deduce their oxidation state once all other oxidation states have been assigned.

Identifying Redox Reactions

• Oxidation: An increase in oxidation state
• Reduction: A decrease in oxidation state

Redox Reactions

• Oxidation and reduction must occur simultaneously.
  • If an atom loses electrons another atom must take them.
• The reactant that causes reduction in another reactant is called the reducing agent.
  • The reducing agent contains the element that is oxidized.
• The reactant that causes oxidation in another reactant is called the oxidizing agent.
  • The oxidizing agent contains the element that is reduced.

Guidelines for Oxidizing and Reducing Agents

Redox reactions:

• Any reaction in which there is a change in the oxidation states of atoms in going from reactants to products.

In a redox reaction:

• The oxidizing agent oxidizes another substance (and is itself reduced).
• The reducing agent reduces another substance (and is itself oxidized).

Activity Series to Determine Spontaneity

• Activity series of metals is a table listing metals in order of decreasing tendency to lose electrons.
  • Metals at the top have higher tendency to undergo oxidation, most reactive.
  • Metals at the bottom have the least tendency to undergo oxidation, least reactive.
• Each reaction in the activity series is an oxidation half-reaction.
• The half-reactions at the top are most likely to occur in the forward direction.
• The half-reactions at the bottom are most likely to occur in the reverse direction.