Chapter 5 Hines

Introduction to Solutions and Aqueous Reactions

Table salt and water form a homogeneous mixture.

Homogeneous mixtures are called solutions. These can be any state of matter: solid, liquid, or gas!

The majority component is the solvent.

The minority component is the solute.

If equal quantities: the one that changes state is the solute. If neither change state, then either can be either.

A solution in which water is the solvent is an aqueous solution.

Solution Concentration

Solutions are often described qualitatively, as dilute or concentrated.

• Dilute solutions have a smaller amount of solute.

• Concentrated solutions have a larger amount of solute.

The above terms usually have the same amount of meaning as “more” and “less”. They are relative and not quantitative.

• We quantify the amount of solute relative to solvent, or concentration of solution.

Solution Concentration: Molarity

A common way to express solution concentration is molarity (M).

Molarity is the amount of solute (in moles) divided by the volume of solution (in liters).

amount of solute (in mol) Molarity (M) volume of solution (in L)

Using Molarity in Calculations

We can use the molarity of a solution as a conversion factor between moles of the solute and liters of the solution.

• For example, a 0.500 Molarity NaCl solution contains 0.500 mol NaCl for every liter of solution.

Solution Dilution

Often, solutions are 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.

  • (M1)(V1) = (M2)(V2)

• moles solute in solution1 = moles solute in solution2

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.

• Volume A → Amount A (moles) → Amount B (moles → Volume B

Charge Distribution in a Water Molecule

There is an uneven distribution of electrons within the water molecule.

This causes the oxygen side of the molecule to have a partial negative charge (δ−) and the hydrogen side to have a partial positive charge (δ+).

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

Solutions containing ions will conduct electricity. These ions are called electrolytes. A solution of salt (an electrolyte) conducts electrical current.

Solution with no ions will not conduct electricity. These solutes are called nonelectrolytes. A solution of sugar (a nonelectrolyte) does not. Sugar is a nonelectrolyte because sugar is a non-ionizing molecule

Strong and Weak Electrolytes

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, or 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 ions when dissolved in water.

• Acids that completely ionize in water are called strong acids.

  • HCl(aq) → H+(aq) + CL-(aq)

• Acids that do not completely ionize in water are called weak acids.

  • HF(aq) ←→ H+(aq) + F-(aq)

The Solubility of Ionic Compounds

When an ionic compound dissolves in water, the resulting solution contains the component ions dissolved in water, not the intact compound.

• 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.

If we mix solid AgNO3 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?

Depends on many factors so predicting whether a compound will dissolve in water is not easy.

The best way to do it is to conduct experiments. We call this method the empirical method.

We have generated a set or “rules” based on empirical observations. In the past students used to memorize these, but you do not need to; just understand how to use them.

Precipitation Reactions

Precipitation reactions are ones in which a solid forms when we mix two solutions.

Reactions between aqueous solutions of ionic compounds sometimes 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. –

   1. Exchange ions.

      1. (+) ion from one reactant with (−) ion from other

   2. Balance charges of combined ions to get the formula of each product.

3. Determine solubility of each product in water.

   1. Use the solubility rules.

   2. 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.

   1. Remember to change only coefficients, not subscripts.

Complete Ionic Equations

An equation showing the complete neutral formulas for each compound in the aqueous reaction as if they existed as molecules is called a molecular equation

• 2 KOH(aq) + Mg(NO3)2(aq) → 2 KNO3(aq) + Mg(OH)2(s)

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 like this one:

• 2 K+(aq) + 2 OH-(aq) + Mg²+(aq) + 2 NO3-(aq) → 2 K+(aq) + 2 NO3-(aq) + Mg(OH)2(s)

Net Ionic Equations reveal what reaction actually occurs

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

Acid–Base Definitions

Arrhenius Definitions:

• Acid:

  • Substance that produces H+ in aqueous solution

  • HCl(aq) → H+(aq) + Cl-(aq)

• In solution, H+ bonds with water to produce the hydronium ion, H3O+.

• 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(aq) → Na+(aq) + OH-(aq)

Monoprotic acids

HCl → H+ + Cl - Strong Electrolyte, Strong Acid

HNO3 → H+ + NO3- - Strong Electrolyte, Strong Acid

CH3COOH ←→ H+ + CH3COO- - Weak Electrolyte, Weak Acid

Diprotic acids

H2SO4 → H+ + HSO4- - Strong Electrolyte, Strong Acid

HSO4- ←→ H+ + SO4 2- - Weak Electrolyte, Weak Acid

Triprotic acids

H3PO4 ←→ H+ + H2PO4- - Weak Electrolyte, Weak Acid

H2PO4- ←→ H+ + HPO4 2- - Weak Electrolyte, Weak Acid

HPO4 2- ←→ H+ + PO4 3- - Weak Electrolyte, Weak Acid

Neutralization Reaction

Acid + Base → Salt + Water

HCl(aq) + NaOH(aq) → NaCl(aq) + H2O

H+ + OH- → H2O

All Na and Cl are eliminated

Acid–Base Reactions

These reactions are called neutralization reactions because the acid and base neutralize each other’s properties.

• 2 HNO3(aq) + Ca(OH)2(aq) → Ca(NO3)2(aq) + 2 H2O(l)

The net ionic equation for an acid–base reaction is:

• H+(aq) + OH-(aq) → H2O

• 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, H3O+.

  • Most chemists use H+ and H3O+ interchangeably.

Bases dissociate in water to form OH− ions.

• Bases, such as NH3, that do not contain OH− ions, produce OH− by pulling H+ 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.

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. BUT THIS IS NOT THE DEFINITION!

Oxidation

loss of electrons

Reduction

gain of electrons

Redox Reactions

often involve incomplete electron transfer--Transfer of electron “ownership”

Electron transfer does not need to be a complete transfer for the reaction to qualify as oxidation–reduction. Example:

• H2(g) + Cl2(g) → 2 HCl(g)

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.

Combustions are Redox Rxns.

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 magnitude before charge: +1 or −1.

Ion charges are real, measurable charges.

• They are written with charge before magnitude: 1+ or 1−

Rules for Assigning Oxidation States

The following rules are in order of priority:

1. Free elements have an oxidation state = 0

• Cu = 0, Cl2 = 0

2. Monatomic ions have an oxidation state equal to their charge.

• Ca2+ = +2, Cl- = -1

3. The sum of the oxidation states of all the atoms in a compound is 0.

• For H2O : H = +1 and O = -2 : 2(+1) + 1(-2) = 0

4. The sum of the oxidation states of all the atoms in a polyatomic ion equals the charge on the ion.

• NO3- : N = +5 and O = -2 : (+5) + 3(-2) = -1

5. Group 1 Metals have an oxidation state of +1 in all of their compounds

• Na = +1 in NaCl

6. Group 2 Metals have an oxidation state of +2 in all of their compounds

• Mg = _2 in MgCl2

Nonmetal Oxidation States

Fluorine = −1

Hydrogen = +1

Oxygen = −2

Group 7A = −1

Group 6A = −2

Group 5A = −3

Extra Oxidation State Rules

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 assigning oxidation states to elements that are not covered by other rules, 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 (loss of electrons)

Reduction: A decrease in oxidation state (gain of electrons)

• C + 2S = CS2

Carbon changes from an oxidation state of 0 to an oxidation state of +4. Carbon loses electrons and is oxidized.

Sulfur changes from an oxidation state of 0 to an oxidation state of −2. Sulfur gains electrons and is reduced.

Oxidizing and reducing agents

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 and reduces the other substance.

The reactant that causes oxidation in another reactant is called the oxidizing agent. The oxidizing agent contains the element that is reduced and oxidizes the other substance.

• 2 Na(s) + Cl2(g) → 2Na+Cl-(s)

• Na is oxidized, while Cl is reduced.

• Na is the reducing agent, and Cl2 is the oxidizing agent.