chapter 13 acids, bases and salts malone and dolter- basic concepts of chemistry 9e2 setting the...

Chapter 13Chapter 13

Acids, Bases and SaltsAcids, Bases and Salts

Malone and Dolter- Basic Concepts of Chemistry 9e

2

Setting the Stage – Varying Properties of Acids

Acids are often the things in foods that give them a distinctive taste (lemon, vinegar) that is generally sour.

Acids are a group of compounds that have similar chemical properties (to a greater or lesser extent).

Acids are also corrosive, as in acid rain: some are dangerously corrosive.


3

Setting a Goal – Part AAcids, Bases, and the Formation of Salts

You will expand your knowledge of acids and bases with more general definitions that allow us to examine their comparative strengths.


4

Objective for Section 13-1

List the general properties of acids and bases.


5

Some Strangely Named Acids

Angelic acid. A gift from heaven? Erotic acid (actually an error, its correct name is

orotic acid). Does it beat ginseng? Moronic acid. Not a very clever name? Tiglic acid. Does it make you roll around the floor,

laughing?


6

13-1 Properties of Acids and Bases

Acids Taste sour (although we don’t taste lab

chemicals!!). React with certain metals (Zn, Fe) with

the liberation of H2 gas. Cause certain organic dyes to change

color.


7

Properties of Acids and Bases, Con’t

Acids React with limestone (CaCO3) with the

liberation of CO2 gas React with bases to form salts and

water Bases

Taste bitter Feel slippery or soapy


8

Properties of Acids and Bases, Con’t

Bases React with oil and grease. Cause certain dyes to change color. React with acids to form salts and

water. Are also known as caustic or alkaline.


9

Composition of Acids and Bases

Arrhenius definition Acids produce H+ in water. Bases produce OH- in water.

E.g.HCl(aq) + H2O(l) H3O+(aq) + Cl-(aq)

Na+OH-(s) + aq Na+(aq) + OH-(aq)


10

The Proton in Water

The proton is represented in water in a variety of ways, but they are all the same thing.

H+(aq) is a common representation. Hydronium, H3O+(aq) tries to indicate that

the solvated proton is complex (really H5O2

+, H7O3+, or similar).

Remember that H+(aq) is not a bare proton in water.


11

Strong Acids in Water

Strong acids completely dissociate into a proton and an anion.

HCl(g) + H2O(l) → H3O+(aq) + Cl-(aq)

Weak acids are only partially dissociated.

HCOOH(aq) + H2O(l) H3O+(aq) + HCOO-(aq)


12

Reactions of Acids

1. Acids react with metals and release H2

Zn(s) + 2 HCl(aq) ZnCl2 (aq) + H2 (g)Net ionic equation: Zn(s) + 2 H+(aq) Zn2+(aq) + H2 (g)

2. Acids react with limestone to release CO2

CaCO3 (s) + 2 HNO3 (aq) Ca(NO3)2 (aq) + H2O(l) + CO2 (g)


13

Reactions of Acids

Net ionic equation: CaCO3 (s) + 2 H+

(aq) Ca2+ (aq) + H2O(l)

+ CO2 (g)

3. Acids react with basesAcid + Base Salt + WaterHClO4 (aq) + NaOH(aq) NaClO4 (aq) + H2O(l)Net ionic equation: H+ (aq) + OH- (aq) H2O(l)


14

Strong Bases in Water

Bases are compounds that produce OH- ions in water.

Such solutions are basic, sometimes referred to as alkaline or caustic solutions.

Most strong bases are salts containing OH-

NaOH(s) → Na+(aq) + OH-(aq)


15

Common Bases

NaOH - caustic soda or lye KOH - caustic potash Ca(OH)2 - slaked lime

NH3 - ammonia

Ammonia is a weak base:

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


16

Objectives for Section 13-2

Identify the Brønsted acids and bases in a proton exchange reaction.

Identify conjugate acid-base pairs.


17

13-2 Brønsted-Lowry Acids and Bases

The Arrhenius definition focuses on molecular compounds and does not account for the basicity of NH3 and CO3

2- .

Acid – a proton donor Base – a proton acceptor For HCl(aq) + H2O(l) H3O+(aq) + Cl-(aq)

HCl is the acid, H2O is the base


18

Conjugate Acid-Base Pairs

Conjugate acid-base pairs are species that differ by a proton.e.g. HCl/Cl-; H3O+/H2O

The conjugate base is the base that remains when an acid donates a proton.

The conjugate acid is the acid that is formed when the base accepts a proton.


19

Amphiprotic Substances

Amphiprotic – species that can both accept and donate protons. Water is the most important example.Water as an acid:

Water as a base:HCl(aq) + H2O(l) → H3O+(aq) + Cl-(aq)

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


20


Calculate the hydronium ion concentration in a solution of a strong acid and a weak acid given the initial concentration of the acid and the percent ionization of the weak acid.


21

13-3 Strengths of Acids and Bases

Related to the concentration of active ingredient (H+ or OH-) produced by the species.

Strong acids dissociate 100% into ions in solution.

There are only six common strong acids:HCl, HBr, HI, H2SO4, HNO3, HClO4


22

Strengths of Acids and Bases

Strength with respect to what? Generally, the strongest acid reacts with

the strongest base to produce a weaker conjugate acid and weaker conjugate base.

In the case of a strong acid in water, the molecular acid (e.g. HCl) is a stronger proton donor than H3O+.


23

Strengths of Acids and Bases Most acids are weak and therefore only partially

ionized. This represents an equilibrium, where a reaction

does not proceed to products completely and is indicated by the double headed arrow.

HA(aq) H+(aq) + A-(aq)

OR

HA(aq) + H2O(l) H3O+(aq) + A-(aq)


24

Equilibrium

The double headed arrow indicates that two reactions are going on: the forward reaction and the reverse, as shown for HF(aq):

ForwardHF(aq) + H2O(l) H3O+(aq) + F-(aq)

ReverseH3O+ (aq) + F- (aq) HF(aq) + H2O(l)


25

Equilibrium At equilibrium, the rates of the forward

and reverse reactions are the same, so it appears that the reaction has stopped since the concentrations no longer change.

It is important to remember that the forward and reverse reactions are continually occurring.

For weak acids, the equilibrium favors the reactants (the molecular species present).


26

% Ionization of an Acid [ ] indicate concentration in moles per liter Concentration times % ionization = concentration

of each ion Consider a 0.20 M solution of a 5% ionized acid:What is its hydrogen ion concentration?

Solution

0.20 M x5.00%

100%= 0.010M


27

Ionizable Protons

Not all protons in an acid are ionizable. Polar bonds can be ionized, but nonpolar

ones cannot, as shown for acetic acid.

C

H

H

H

C

O

O H

Polar, ionizablebondNonpolar

bonds, notionizable


28

The Strength of Bases

All alkali metal hydroxides are strong and very water soluble.

Alkaline earth hydroxides (except Be(OH)2) also completely dissociate into ions.

The exception is Mg(OH)2, which is not very water soluble, hence it produces few ions in water (very little OH-).


29


Write the molecular, total ionic, and net ionic equation for neutralization reactions.


30

13-4 Neutralization and the Formation of Salts

Neutralization is a double replacement reaction.

The active ingredient from the acid (H+) reacts with the active ingredient from the base (OH-) to form the molecular compound water.

A salt is what is left over – usually as spectator ions if the salt is soluble.


31

Total and Net Ionic Equations for Neutralization

Total ionic equationH+(aq) + Cl-(aq) + Na+(aq) + OH-(aq)

→ Cl-(aq) + Na+(aq) + H2O(l) Net ionic equation

H+(aq) + OH- (aq) → H2O(l)


32

Three Kinds of Acid-Base Reactions

1. 2 HNO3 (aq) + Ca(OH)2 (aq) Ca(NO3)2 (aq) + 2 H2O (l)strong acid plus strong base, but Ca(OH)2 produces two OH-, so we need two moles of acid per mole of base.


33

Three Acid-Base Reactions

2. HClO(aq) + LiOH(aq) LiClO(aq) + H2O(l)weak acid plus strong base. Since acid is weak, it is represented as its molecular form:HClO(aq) + OH- (aq) ClO- (aq) + H2O(l)


34

Three Acid-Base Reactions

3. H2SO4 (aq) + 2 NaOH (aq) Na2SO4 (aq) + 2 H2O (l)

First ionization of H2SO4 is complete but the second is not.e.g.H2SO4 (aq) HSO4

- (aq) + H3O+ (aq) complete

HSO4- (aq) SO4

2- (aq) + H3O+ (aq) not complete

One mole of NaOH with one mole of H2SO4 yields a partial neutralization forming water and NaHSO4 (sodium bisulfate), an acid salt.


35

Acid Salts

Acid salts are ionic compounds containing an anion with one or more acidic hydrogens that can be neutralized by a base.

Can be named two ways:NaHSO4

sodium bisulfate or sodium hydrogen sulfate.


36

Polyprotic Acids

H2SO4 is an example of a polyprotic acid, one that has more than one ionizable proton.

Monoprotic HCl Diprotic H2SO4

Triprotic H3PO4


37

Setting a Goal – Part BThe Measurement of Acid Strength

You will learn about how the relative acidities of aqueous solutions are expressed.


38


Using the ion product of water, relate the hydroxide ion and the hydronium ion concentrations.


39

13-5 Equilibrium of Water

Acidity and basicity of compounds are compared to those of water.

Water undergoes autoionization (note: reaction lies far to the left).


40

Autoionization of Water

From experiment, at 25ºC [H3O+] = [OH-] = 1.0 × 10-7

M

Kw = [H3O+] [OH-] = 1.0 × 10-14 M

Kw is called the ion product of water

Neutral solution[H3O+] = [OH-]

Acidic solution [H3O+] > [OH-]

Basic solution [H3O+] < [OH-]


41

The Relationship Between H3O+ and OH– in Water


42


Given the hydronium ion concentration, calculate pH and pOH.


43

13-6 The pH Scale

Hydronium and hydroxide concentrations are small numbers (~10-7 mol/L), so a notation was developed to abbreviate it:pH = -log[H3O+]pOH = -log[OH-]

Neutral pH = pOH = 7 Acidic pH < 7; pOH > 7 Basic pH > 7; pOH < 7


44

Significant Figures and Logs

Note that the number of digits to the right of the decimal point in the result of a log indicates the number of significant digits.

If [H3O+] = 1.00 × 10-4 M, then pH = 4.000, where the three zeroes to the right of the decimal are the three significant digits from the hydronium ion concentration.


45

pH and pOH

Since pH = -log[H+], pOH = -log [OH-], and

pKW = -logKW = 14, there is a simple

relationship between pH and pOH:

pH + pOH = 14.00


46

pH and pOH


47

The pH Scale


48

Setting a Goal - Part CSalts and Oxides as Acids and Bases

You will learn how salts and oxides can also affect the pH of aqueous solutions


49


Identify a salt solution as neutral, acidic, or basic.


50

13-7 Effects of Salts on pH - Hydrolysis

Hydrolysis is literally cleavage by water. In this case, it is the reaction of a cation as

an acid or an anion as a base. Anions of strong acids do not undergo

hydrolysis, but those of weak acids do.


51

Anion Hydrolysis

Anions of weak acids undergo hydrolysis to yield molecular acid plus hydroxide.ClO-(aq) + H2O(l) HClO(aq) + OH-(aq)

Generally, the weaker the acid, the stronger the conjugate base (and vice versa).


52

Cation Hydrolysis - Ammonia

Ammonia is a weak base, hence the reaction:NH3(aq) + H2O(l) NH4

+(aq) + OH-(aq)does not proceed very far

Rather, the reverse type of reaction predominates: NH4

+(aq) + H2O(l) H3O+(aq) + NH3(aq) The ammonium ion undergoes hydrolysis

to yield molecular base plus hydronium ion


53

Salts in Water – Effect on pH

1. Neutral solutions of salts – neither the cation nor the anion undergoes a hydrolysis reaction. Therefore pH stays the same. Examples are NaCl, CsNO3

2. Basic solutions of salts – salt from the cation of a strong base and the anion of a weak acid. Cation does not affect pH but the anion reacts with water to produce hydroxide. Examples are NaOCl, NaF, CH3COONa


54

Salts in Water – Effect on pH

3. Acidic solutions of salts – cation of a weak base and the anion of a strong acid. The anion does not affect the pH, but the cation reacts to produce H+. Primary example is NH4

+. Also N2H5+ and

Al3+

4. Complex cases – examples that have to be measured in order to understand. Consider NaHCO3 or CH3COONH4


55

Objectives for Section 13-8

Describe how a buffer solution works.

Write equations illustrating the fate of hydronium or hydroxide ions added to such a solution.


56

13-8 Control of pH – Buffer Solutions

Buffer solutions resist changes in pH caused by the addition of a limited amount of a strong acid or a strong base.

Requires two species; one that can react with added acid and one that can react with added base.


57

The Composition of a Buffer

Generally a buffer can be made from: A weak acid and the salt of its conjugate

base: CH3COOH/CH3COO- is an example A weak base and the salt of its conjugate

acid: NH4+/NH3 constitutes an example


58

Control of pH - Buffers

A common buffer system (also found in blood)–NaH2PO4/Na2HPO4

Buffering action: Added acid

HPO42- + H3O+ H2PO4

- + H2O Added base

H2PO4- + OH- HPO4

2- + H2O


59

Buffer Capacity

The amount of acid or base that can be added to a given amount of a buffer solution without the pH changing beyond a given limit.


60


Determine, if possible, whether a specific oxide is acidic or basic.


61

13-9 Oxides as Acids and Bases

An ecologically important example is SO2

Combustion of coal that contains sulfur also produces SO2

In the atmosphere:2 SO2(g) + O2(g) 2 SO3(g) SO3(g) + H2O(g) H2SO4(aq)

SO3 is an example of an acid anhydride (an acid without water).


62

Anhydrides

Many nonmetal oxides are acid anhydrides- CO2, for example

Base anhydrides are ionic metal oxides that dissolve in water to form bases.Na2O(s) + H2O(l) 2 NaOH(aq)

Salts can be formed by reacting an acid anhydride with a base anhydride.SO2(g) + CaO(s) CaSO3(s)

chapter 13 acids, bases and salts malone and dolter- basic concepts of chemistry 9e2 setting the...

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