Acids, Bases, and Salts Chapter 23 Properties of Acids Acid Property #1. The word acid comes from...

Acids, Bases, and Salts

Chapter 23

Properties of Acids

Acid Property #1. The word acid comes from the Latin word acere, which means "sour." All acids taste sour.

Acid Property #2. In 1663, Robert Boyle wrote that acids would make a blue vegetable dye called "litmus" turn red.

Acid Property #3. Acids destroy the chemical properties of bases.

Acid Property #4. Acids conduct an electric current. (electrolytes)

Acid Property #5. Upon chemically reacting with an active metal, acids will evolve hydrogen gas (H2).

Properties of Bases

Base Property #1. The word "base" has a more complex history and its name is not related to taste. All bases taste bitter.

Base Property #2. Bases are substances which will restore the original blue color of litmus after having been reddened by an acid.

Base Property #3. Bases destroy the chemical properties of acids.

Base Property #4. Bases conduct an electric current. (electrolytes)

Base Property #5. Bases feel slippery, sometimes people say soapy. This is because they dissolve the fatty acids and oils from your skin and this cuts down on the friction between your fingers as you rub them together.

Properties of A Salt

A salt is the combination of a cation(+ ion) and an anion (- ion).

Salts are products of the reaction between acids and bases.

Solid salts usually make crystals. If a salt dissolves in water solution, it

usually dissociates into the anions and cations that make up the salt.

The Acid Base Theory

The three main theories regarding acids and bases are:

1. Arrhenius 3. Lewis

2. Brønsted-Lowry

Arrhenius Theory – late 1890s

Acid - any substance which delivers hydrogen ion (H+) to the solution.

HA ---> H+ + A¯ Base - any substance which delivers

hydroxide ion (OH¯) to the solution.XOH ---> X+ + OH¯

When acids and bases react, they neutralize each other, forming water and a salt:

HA + XOH ---> H2O + XA

Problems with Arrhenius Theory

The theory did not explain why ammonia (NH3) was a base.

The solvent has no role to play in this theory. We know however, that an acid added to benzene will not dissociate. Solvents are vital.

The end result of mixing certain acids and bases can be a slightly acidic or basic solution. Arrhenius had no explanation for this phenomenon.

Brønsted – Lowry Theory – Early 1920s

Two chemists, independent of one another, proposed a new definition of an acid and a base.

An acid is a substance from which can donate a proton.

A base is a substance that can accept a proton from an acid.

*Can occur in any solvent.

Reactions Based on Bronsted - Lowry

Reactions that proceed to a large extent: HCl + H2O <===> H3O+ + Cl¯ HCl - this is an acid, because it has a proton

available to be transferred. H2O - this is a base, since it gets the proton

that the acid lost. Now, here comes an interesting idea: H3O+ - this is an acid, because it can give a

proton. (hydronium) Cl¯ - this is a base, since it has the capacity

to receive a proton.

HCl + H2O <===> H3O+ + Cl¯

Notice that each pair (HCl and Cl¯ as well as H2O and H3O+ differ by one proton (symbol = H+). These pairs are called conjugate pairs.

HNO3 + H2O <===> H3O+ + NO3¯

The acids are HNO3 and H3O+ and the bases are H2O and NO3¯.

Bases and Conjugate Acid

Base NameConjugate acid Name

CH3OO-

Acetate ion CH3COOH Acetic acid

NH3 Ammonia NH4+ Ammonium

H2PO4-

Dihydrogen phosphate ion

H3PO4 Phosphoric acid

HSO4- Hydrogen sulfate

ionH2SO4 Sulfuric

acid

OH- Hydroxide ion H20 water

NO3- Nitrate ion HNO3 Nitric acid

H2O water H30+ Hydronium ion

Weak/Strong Conjugate Bases

The stronger the acid the weaker the conjugate base:HCl (SA) Cl- (weak conjugate base)

The weaker the acid the stronger the conjugate base:H3PO4 (WA) H2PO4

- (strong conjugate base)

Water is considered an amphoteric substance because it can act as either an acid or a base.

EX: HSO4-1, HPO4

-2

Lewis Theory –Early 1920s Remember drawing Lewis Dot Structures

for ionic and covalent compounds? Lewis Theory focuses on the nature of

electrons rather than proton transfer. An acid as an electron pair acceptor and

a base as an electron pair donor. Lewis Theory is much more general and

apply to reactions that do not involve hydrogen or hydrogen ions.

Draw the Lewis Dot Structure for each molecule.

The molecule with a lone e- pair on the central atom is the Lewis base. The other is the Lewis acid. It will accept the e- pair.

Identify each as a Lewis Acid or Lewis Base:

1. NH3 7. Fe+2

2. PCl33. H2O

4. AlCl35. SO3

6. O-2

Strong Acids and Bases

Strong acids are those that ionized completely in water.

The dissociation of a strong base looks like the diagram at the right in that it dissociates into positive and negative ions.

Acids

1. Monoprotic- one hydrogen to donateEX: HCl

2. Polyprotic- two to three hydrogens to donate. EX: diprotic: H2SO4

triprotic: H3PO4

*H+ exists as H3O+ in aqueous solutions.

Weak Acids and Bases

Some acids and bases ionize only slightly in water.

These are considered weak.

The most important weak base is ammonia.

Acidic NeutralSolution Solution

Strong Acids

*HNO3 - nitric acid *HCl - hydrochloric acid*H2SO4- sulfuric acid

*HClO3 - chloric acid

*HClO4 - perchloric acid *HBr - hydrobromic acid*HI - hydroiodic acid

Strong Bases

*LiOH - lithium hydroxide*NaOH - sodium hydroxide*KOH - potassium hydroxide*RbOH - rubidium hydroxide*CsOH - cesium hydroxide*Ca(OH)2 - calcium hydroxide*Sr(OH)2 - strontium hydroxide*Ba(OH)2 - barium hydroxide

Neutralization Reactions

The word "neutralization" is used to describe the reaction of an acid plus a base because the acid and base properties of H+ and OH- are destroyed or neutralized.

In the reaction, H+ and OH- combine to form water and a salt.

When acids and bases are equal in strength and concentration, a neutral (pH = 7) solution is formed.

A neutralization reaction is a type of double replacement reaction.

EX:a. HCl + NaOH --> b. H2SO4 + Fe(OH)3 -->

Calculations of Neutralization Reactions

Utilyze equation: MAVA = MBVB MA and VA= Molarity and Volume of Acid MB and VB= Molarity and Volume of Base

*Volumes may remain in milliliters*Must adjust the Molarity of all acids or

bases based on the number of moles of H+ or OH- they contribute. (See next slide)

*Add

Adjust MA of H2SO4 by 2 (times 2 moles H+). Only acid that needs adjusted.

Adjust MB by 2 for the soluble group 2 hydroxides: Ca, Sr, Ba.

If finding Molarity of these substances, divide molarity answer by 2 to get actual molarity. (opposite process)

1. How many mL of a 1.5 M HCl acid is needed to neutralize a 500. mL 1.5 M NaOH solution?

2. How many mL of a 2.0 M H2SO4 acid is needed to neutralize a 500. mL 1.5 M KOH solution?

3. How many mL of a 0.750 M HNO3 acid is needed to neutralize a 275 mL 1.5 M Ca(OH)2 solution?

4. Calculate concentration of acid if 300. mL of H2SO4 is used to neutralize a 500. mL 2.50 M NaOH solution.

Net Ionic Equations For aqueous reactions, it is common to

write equations in the ionic form.Standard form: NaOH + HCl NaCl + H2OIonic form: Na+ + OH- + H+ + Cl- Na+ + Cl- + H2ONotice* substances occurring in molecular

form (H2O) are written as molecules.

* ionic substances are written as ions if soluble.

Now take the ionic equation and cancel out any spectator ions on the product and reactant side of the equation.

The result will be a net ionic equation.Na+ + OH- + H+ + Cl- Na+ + Cl- + H2O

OH- + H+ H2O

Rules for Writing Net Ionic Equations

Rule 1 Binary Acids: HCl, HBr, and HI are strong: all other binary acids and HCN are weak. Strong acids are written in ionic form; weak acids are written in molecular form.

Rule 2 Ternary Acids: If the number of oxygen atoms in an inorganic acid molecule exceeds the number of hydrogen atoms by two or more, the acid is strong. We will consider all organic carboxylic acids as weak.

Strong: HClO3, HClO4, H2SO4, HNO3, H2SeO4

Weak: HClO, H3AsO4, H2CO3, H4SiO4, HNO2

Rule 3 Polyprotic Acids: (acids that contain more than one ionizable hydrogen atom. EX: H2SO4, H3PO4, H2CO3). In the second and subsequent ionizations H2SO4 is always weak, even though it is a strong acid.

Rule 4 Bases: Hydroxides of Group 1 and 2 (Ca, Sr, and Ba) are soluble and strong. All others including ammonia, hydroxlamine, and organic bases are weak.

Rule 5 Salts: Salts are written in ionic form if soluble, and in undissociated form if insoluble. *Know the solubility rules.

Rule 6 Oxides: Oxides (except Group I metals) are always written in molecular or undissociated form. MgO, ZnO (H2O)

Rule 7 Gases: Gases are always written in molecular form. CO2, NH3, O2,

Practice Net Ionic Equations

1. AgNO3 (aq) + H2SO4 (aq)

2. H4SiO4 (aq) + NaOH (aq)

3. HBr (aq) + KOH (aq) 4. HCl(aq) + Cr(NO3)2(aq) + HgCl2(aq)

CrCl3(aq) + Hg2Cl2(s) + HNO3(aq)

5. H2CO3(aq) + NaOH(aq)

pH and pOH

The scale is measured on a log scale of 0 to 14, with each unit representing a ten-fold change.

pH Scale The pH scale is a measure of hydronium ion

[H3O+] concentration (acid molarity) as well as the hydroxide [OH-] (base molarity) concentration.

Hydronium ion concentration indicates acidity and will have a pH of 0-6.9 .

Hydroxide ion concentration indicates basicity and will have a pH of 7.1-14 .

The higher the [H3O+], the higher the acidity, the higher the [OH-], the higher the basicity.

Calculating pH

The concentration (M or mol/L) of H3O+ is expressed in powers of 10, from 10-14 to 100.

Scientists use pH which is the negative log of acid concentration, [H3O+], and can be a negative value if molarity is greater than 1.

pH = -log[H3O+] acid molarity

[H+]

EX: 0.50M HCl is added to water to make a final volume of 1 liter. What is the pH of this solution?

Step 1: Identify acid concentration, [H3O+], in mol/L

0.50 MStep 2: Place value in equation and

solve. pH = -log[0.50] = 0.30 acidic

Practice pH Calculations

Find pH of the following solutions if [H3O+] is:

1. 1.00 x 10-3

2. 6.59 x 10-6

3. 9.47 x 10-10

Calculate pH of Strong Acids

If you have the strong polyprotic acid, H2SO4, you must adjust the molarity by the number of moles of H+ contributed (2 moles of H+).

EX: 0.250M H2SO4 is added to water to make a final volume of 1 liter. What is the pH of this solution?

Step 1: Identify [H3O+] in mol/L0.250 M

Step 2: Recognize that H2SO4 contributes 2 moles H+.Step 3: Place value in equation and solve.

pH = -log[0.250 x 2] = 0.301 acidic

Calculate pH of Weak Acids

Weak acids will not dissociate 100%. A dissociation factor will be included in these problems.

EX: Calculate pH of a 0.150 M HNO2 solution (dissociation is 5.00%).

Step 1: 0.150 MStep 2: Convert % to decimal: 5.00%

0.0500Step 3: pH = -log[0.150 x 0.0500] = 2.13

acidic

Calculate H30+ and H+ Concentration from pH

If given pH you can calculate the hydronium or hydrogen ion (acid) concentration by performing the anti-log function.

EX: pH is 2.00. 10^(-2.00) = 1.00 x 10-2M Find [H3O+] if the pH is:

1.6.678 3. 10.02.2.533 4. 2.56 Remember the unit for concentration is M.

pOH

You can calculate the pH of a solution if you know the concentration of hydroxide ion [OH-] (base molarity).

If we use the ion product constant of water (1.00 x 10-14)we can derive this equation:

[pH]x[pOH] = 1.00 x 10-14

Working with this equation leads to:pH + pOH = 14

EX: Find the pH of a solution with an [OH-] of 1.0 x 10-8 M.

Step 1: Calculate pOH by using equation: pOH = -log[OH-]

pOH = -log[1.0x10-8] = 8.0Step 2: Subtract the pOH from 14 to find

pH:14 – pOH = pH

14 – 8.0 = 6.0 acidic

EX: Find pH of a 0.750 M KOH solution.

Step 1: Identify that your substance is a base and you need to calculate pOH.pOH = -log[0.750] = 0.125

Step 2: 14 – 0.125 = 13.875 pH basic

Practice pH Calculations Using pOH

Find the pH of the following solutions with [OH] of:

1. 1.00 x 10-4M2. 2.64 x 10-13M3. 5.67 x 10-2M4. 3.45 x 10-11M

Calculate [OH-] from pH or pOH

If given pH:14 – pH = pOH10^(-pOH) = [OH-]

If given pOH:10^(-pOH) = [OH-]

Remember unit for concentration is M.

Summary of pH and pOH

pH = -log[H+] or [H3O+] (acid molarity)

pOH = -log[OH-] (base molarity)pH + pOH = 14[H+ or H3O+] (acid molarity) = 10^(-pH)

[OH-] (base molarity) = 10^(-pOH)*Hints:1. Identify initial substance as acid or base.2. Label all concentrations with M unit.3. Adjust molarity of strong and weak acids/bases.

To check an answer for concentration, use equation [OH-][H3O+] = 1.00 x 10-14

EX: A problem gives [OH-] as 2.30 x 10-

4M. You find the pOH to be 3.64 and pH is 10.36. You find [H3O+] by 10^(-10.36) to be 4.365 x 10-11M.

Check answer by dividing 1.00 x 10-14 by the given [OH-] 2.30 x 10-4. You get 4.35 x 10-11. This answer is comparable to the original and is acceptable.

pH Indicators

acid-base indicator:A substance that indicates the degree of acidity or basicity of a solution through characteristic color changes.

The narrower the pH range, the better the indicator.

Refer to the following chart to identify common indicators and their pH ranges.

Titrations

Titration is a standard laboratory method of quantitative/chemical analysis which can be used to determine the concentration of an unknown reactant (acid or base).

An acid or base of known concentration (a standard solution) and volume is used to react with a measured volume of an unknown concentration of an acid or base.

Using a buret to add the [unknown], it is possible to determine the exact amount (V) that has been consumed when the endpoint is reached.

The endpoint is the point at which the titration is stopped.

This is classically a point at which the number of moles of [unknown] is equal to the number of moles of [known].

Many methods can be used to indicate the endpoint of a reaction; titrations often use visual indicators.

In simple acid-base titrations a pH indicator may be used, such as phenolphthalein, which turns (and stays) pink when a certain pH (pH > 7) is reached or exceeded.

Methyl orange can also be used, which is red in acids and yellow in alkalis (bases).

Any indicator that changes color along the steep portion of the titration curve is suitable for the titration. Methyl violet changes color too soon, and alizarin yellow

R too late.

Due to the logarithmic nature of the pH curve, the transitions are generally extremely sharp, and thus a single drop of unknown just before the endpoint can change the pH by several points - leading to an immediate color change in a chosen indicator.

Titrations of Strong Acids and Strong Bases

Titration of a strong acid and a strong base will result in an equivalence point of 7 because a neutral salt water solution is formed.

Titration of a Strong Acid and a Weak Base

Titration of a strong acid and a weak base will result in an equivalence point of less than 7 because an acidic salt water solution is formed.

Titration of a Weak Acid and a Strong Base

Titration of a weak acid and a strong base will result in an equivalence point of greater than 7 because a basic salt water solution is formed.

Titrations of polyprotic acids result in titration curves with more than one equivalence point.

The titration to the right is that of H3PO4

Titrations of weak acids and weak bases require calculations and result in titration curves without a sharp transition.