chapter 10: acids and bases

Chapter 10: Acids and Bases

• When we mix aqueous solutions of ionic salts, we are not mixing single components, but rather a mixture of the ions in the solid– The ionic solid dissolves in the water

• We call a compound that dissolves in water soluble and if it doesn’t, it is insoluble

Electrolytes

• When an ionic compound dissolves in water, it forms an electrolyte solution– The compound may be a strong

electrolyte if it dissolves completely or a weak electrolyte if it only partially dissolves (doesn’t exist entirely as ions in solution)

Precipitation Reactions• A precipitation reaction takes place

when solutions of 2 strong electrolyte solutions are mixed and react to form an insoluble solid

Complete and Net Ionic Equations

AgNO3 (aq) + NaCl (aq) --> AgCl (s) + NaNO3 (aq)

• A Complete Ionic Equation shows all of the ions and solids in a precipitation reaction

Complete Ionic Equation:Ag+(aq) + NO3

-(aq) + Na+(aq) + Cl-(aq) --> AgCl(s) + Na+(aq) + NO-3 (aq)

Complete and Net Ionic Equations

• A Net Ionic Equation removes the spectator ions from the complete ionic equation– Spectator Ions don’t do anything in the reaction and are

found on both sides of the arrow.

Complete Ionic Equation:Ag+(aq) + NO3

-(aq) + Na+(aq) + Cl-(aq) --> AgCl(s) + Na+(aq) + NO-3 (aq)

Net Ionic Equation:Ag+(aq) + Cl-(aq) --> AgCl(s)

AgNO3 (aq) + NaCl (aq) --> AgCl (s) + NaNO3 (aq)

Acids and Bases

• There are several possible definitions of acids and bases, but we’ll start with the Bronsted definition initially

A Bronsted Acid is a Proton DonorA Bronsted Base is a Proton Acceptor

• Acids are only acids once they donate their proton to an accepting base

• Bases are only bases once they accept a proton from a donor

HCl and Phase

• In the gas phase, HCl is just another molecule with 2 atoms

• Once we add the molecule to water however…

Strong and Weak Acids

HCl (aq) + H2O (l) --> H3O+ (aq) + Cl-(aq)

• The reaction goes almost to completion (K is very ____), so we only draw a single arrow.– HCl is a strong acid

HCN (aq) + H2O (l) --> H3O+ (aq) + CN-(aq)

• The K value for this reaction is low, so the reaction favors the _______– HCN is a weak acid

A Strong Acid is fully deprotonated in solutionA Weak Acid is only partially deprotonated in

solution

Strong and Weak Bases

• A Bronsted base is a proton acceptor• This means it has a lone pair to accept

the proton (more on this in a little bit…)

• Let’s look at CaO:CaO (aq) + H2O (l) --> Ca(OH)2 (aq)

Ca2+(aq) + O2-(aq) + H2O(l) --> Ca2+(aq) + 2OH-(aq)

O2-(aq) + H2O(l) --> 2OH-(aq)

The K value for this reaction is very high and oxide ions are strong bases in water

Strong and Weak Bases

NH3 (aq) + H2O (l) --> NH4+ (aq) + OH- (aq)

• NH3 is electrically neutral, and it has a lone pair to accept the proton, but the K value for the reaction is very low

• Ammonia is a weak base• All amines, organic derivatives of ammonia, are

weak bases

Conjugate Acids and Bases

• The products of proton transfer may also react with water

HCN (aq) + H2O (l) CN- (aq) + H3O+ (aq)

• The cyanide ion may take/accept a proton to reform HCN– This is called a Conjugate Base

• The HCN formed when CN- accepted a proton is called the Conjugate Acid of CN-

The Conjugate Base of an acid is the species left when the acid donates a proton

The Conjugate Acid is the species formed when the base accepts a proton

Lewis Acids and Lewis Bases

•Because of the sheer possibilities that exist in the chemical world, we need to expand our definition of acids and bases to include more than just protons.

A Lewis Acid is an electron pair acceptor

A Lewis Base is an electron pair donor

Lewis Acids and Bases

We’ll use Lewis structures to show how electron pairs move in the reactions of Lewis acids and bases.

Oxide anion reacting with water •The oxide anion is a Lewis base (electron pair donor)

Ammonia reacting with water •The lone pair in Nitrogen grabs a water proton

Carbon dioxide accepts an electron pair from the oxygen of water

Acidic, Basic and Amphoteric Oxides1. Acidic oxides react with water to form a Bronsted acid

CO2(g) + H2O(l) H2CO3(aq)

1. Acidic oxides are molecular compounds of nonmetal oxides

• Basic oxides react with water to form a Bronsted base

CaO(s) + H2O(l) --> Ca(OH)2(aq)

• Basic oxides are ionic compounds of metals

1. Oxides of the metalloids are amphoteric meaning that they react with both acids and bases

Al2O3(s) + 6HCl(aq) --> 2AlCl3 + 3H2O(l)

Al2O3(s) + 2NaOH(aq) --> 2Na[Al(OH)4](aq)

Autoprotolysis

Water is both an acid and a base

H2O(l) + O2-(g) --> 2OH- (water as an acid)

H2O(l) + HCl(aq) --> H3O+ + OH- (water as a base)

Water is Amphiprotic meaning that it can act as a proton donor or proton acceptor

Autoprotolysis

• Because water is amphiprotic, proton transfer between water molecules spontaneously happens– In fact, water is never just H2O

2H2O(l) H3O+ + OH-

This is autoprotolysis

We can describe K as:

€

K = [H3O

+][OH−]

[H2O]2

K = [H3O+][OH−]

Autoprotolysis

Kw = [H3O+][OH-]

• From experiments, we can measure the concentrations of H3O+ and OH- and find them to be equal and 1.0x10-7 M

Kw = [H3O+][OH-]=(1.0x10-7)(1.0x10-7) = 1.0x10-14

Kw is still an equilibrium constant, so whatever we do to one product, the other will compensate to maintain Kw = 1.0x10-14

The pH Scale

pH = -log[H3O+]

• In a pure water sample, the [H3O+] = 1.0x10-7M and the pH is 7.00

• At values lower than 7, the [H3O+] is increasing

• At values higher than 7, the [H3O+] is decreasing (and the pOH is increasing)