Acids and Bases

ACIDS AND BASES

SOME PROPERTIES OF ACIDS

Produce H+ (H3O+) ions in water. The Hydronium Ion (H3O+) is an H+ (proton)

attached to a water molecule. Taste Sour. React with certain metals to produce H2 gas

and a salt. Salt – ionic-metal or a positive polyatomic ion

bonded with a negative ion other than OH-

Example: MgCl2, NH4Cl

Aqueous solutions of acids conduct electricity. Electrolytes – the greater the concentration of ions

in solution, the greater the electrical conductivity. Strong Acids & weak Acids

SOME PROPERTIES OF ACIDS

React with carbonates and bicarbonates to produce carbon dioxide gas.

HCl(aq) + NaHCO3(aq) NaCl(aq) + H2O(l) + CO2(g)

React with bases to form a salt and water. Neutralization Reaction (Double Replacement)

HCl(aq) + NaOH(aq) NaCl(aq) + H2O(l)

pH is less than 7 pH scale expresses the amount of H+ as a number

from 0 to 14 pH of 0 is strongly acidic and has the highest amount of H+

ions, pH of 7 is neutral, pH of 14 is strongly basic and has the fewest H+ ions.

“ABC easy as 123”

Cause acid-base indicators to change color. Acids turn Blue litmus Red

ACID NOMENCLATURE

ACID NOMENCLATURE

Use the flowchart to name the following acids.

HBr

H2CO3

H2SO3

ACID NOMENCLATURE

ACID NOMENCLATURE


HBr Hydrobromic Acid

H2CO3 Carbonic Acid

H2SO3

ACID NOMENCLATURE


HBr Hydrobromic Acid

H2CO3 Carbonic Acid

H2SO3 Sulfurous Acid

GOING BACKWARDS…

Write H first.

Write the 2nd ion. (you may have to check table E)

Assign charges.

Criss Cross, if necessary.

Examples

Sulfuric Acid _______________________ Nitrous Acid _______________________ Oxalic Acid _______________________

GOING BACKWARDS…SULFURIC ACID

Write H first.

Write the 2nd ion. (you may have to check table E) -ate becomes -ic going backwards -ic becomes -ate -ite becomes -ous going backwards -ous becomes -

ite

Assign charges.


GOING BACKWARDS…

Write H first.


Assign charges.


Examples

Sulfuric Acid _____ H2SO4__________ Nitrous Acid _______________________ Oxalic Acid _______________________

GOING BACKWARDS…

Write H first.


Assign charges.


Examples

Sulfuric Acid _____ H2SO4__________ Nitrous Acid _______HNO2 __________ Oxalic Acid ______ H2C2O4_________

NAME ‘EM

HI(aq)

HCl(aq)

H2SO3

HNO3

HClO4

NAME ‘EM

HI(aq)

hydroiodic Acid

HCl(aq)

hydrochloric Acid

H2SO3

Sulfurous Acid

HNO3

Nitric acid

HClO4

Perchloric Acid

SOME PROPERTIES OF BASES

Produce OH- ions in water. Taste Bitter, chalky. Aqueous solutions of bases conduct

electricity. Electrolytes – the greater the concentration of

ions in solution, the greater the electrical conductivity. Strong bases & weak bases

Feel soapy, slippery. This is because they break down the normal

body fat in your hands or whatever part of your body they come into contact with. NaOHBefor

eAfter

SOME PROPERTIES OF BASES

React with acids to form a salt and water. Neutralization Reaction (Double Replacement)

HCl(aq) + NaOH(aq) NaCl(aq) + H2O(l)

pH is greater than 7 pH scale expresses the amount of H+ as a

number from 0 to 14 pH of 0 is strongly acidic and has the fewest OH- ions,

pH of 7 is neutral, pH of 14 is strongly basic and has the greatest amount of OH- ions.

“ABC is easy as 123”

Cause acid-base indicators to change color. Bases turn Red litmus Blue

NAMING BASES

Name the Metal. If the metal has only one possible charge, just write it’s

name. If the metal has more than one possible charge, use

Roman Numerals to indicate the charge.

Follow with Hydroxide

Examples: LiOH

Fe(OH)3

NAMING BASES

Name the Metal.

If the metal has only one possible charge, just write it’s name.

• Li+1 = Lithium If the metal has more than one possible charge, use

Roman Numerals to indicate the charge.• Fe+2 = Iron(II)• Fe+3 = Iron (III)

Follow with Hydroxide Lithium Hydroxide Iron(III) Hydroxide

NAMING BASES

Name the Metal. If the metal has only one possible charge, just write it’s

name. If the metal has more than one possible charge, use

Roman Numerals to indicate the charge.

Follow with Hydroxide

Examples: LiOH Lithium Hydroxide

Fe(OH)3 Iron(III) Hydroxide

SOME COMMON BASES

NaOH____________________ lye

KOH ____________________ liquid soap

Ba(OH)2 ____________________

Mg(OH)2 ____________________

Al(OH)3 ____________________ Maalox

stabilizerfor

plastics

milkof

magnesia

SOME COMMON BASES

NaOH__Sodium Hydroxide_ lye

KOH Potassium Hydroxide_ liquid soap

Ba(OH)2 Barium(II) Hydroxide_

Mg(OH)2 Magnesium(II) Hydroxide

Al(OH)3 Aluminum(III) Hydroxide Maalox

Table ‘L’ lists some common Bases

stabilizer

for plasticsmilk

ofmagnesi

a

GOING BACKWARDS…

Write the symbol of the metal.

Write OH-

Assign Charges.

Criss Cross, if necessary

Examples Cesium Hydroxide Chromium(III) Hydroxide Strontium Hydroxide

GOING BACKWARDS…CESIUM HYDROXIDE


Write OH

Assign Charges


GOING BACKWARDS…


Write OH-

Assign Charges.


Examples Cesium Hydroxide CsOH Chromium(III) Hydroxide Strontium Hydroxide

GOING BACKWARDS…


Write OH-

Assign Charges.


Examples Cesium Hydroxide CsOH Chromium(III) Hydroxide Cr(OH)3

Strontium Hydroxide SrOH

PRACTICE

HBr

H2SO3

H2C2O4

HClO

Ca(OH)2

AgOH

HgOH

HF

HI

HClO4

HCl

LiOH

Sn(OH)2

Ti(OH)3

Name each of the following…

PRACTICE

HBr hydrobromic acid

H2SO3 sulfurous acid

H2C2O4 oxalic acid

HClO hypochlorous acid

Ca(OH)2 calcium(II) hdroxide

AgOH silver hydroxide

HgOH mercury hydroxide

HF hydroflouric acid

HI hydroiodic acid

HClO4 perchloric acid

HCl hydrochloric acid

LiOH lithium hydroxide

Sn(OH)2 tin(II) hydroxide

Ti(OH)3 titanium(III) hydroxide

Name each of the following…

PRACTICE

nitric acid

carbonic acid

dichromic acid

acetic acid

nitrous acid

potassium hydroxide

cesium hydroxide

barium(II) hydroxide

aluminum(III) hydroxide

strontium(III) hydroxide

Now go backwards….

PRACTICE

nitric acid HNO3

carbonic acid H2CO3

dichromic acid H2Cr2O7

acetic acid HC2H3O2

nitrous acid HNO2

potassium hydroxide KOH

cesium hydroxide CsOH

barium(II) hydroxide Ba(OH)2

aluminum(III) hydroxide

Al(OH)3

strontium(III) hydroxide

Sr(OH)3

Now go backwards….

EXPLAINING ACIDS AND BASES

There have been several attempts to explain the properties of acids and bases.

These explanations define how acids and bases behave.

There are three such definitions. Arrhenius Theory Brønsted – Lowry Lewis Acids & Bases

ARRHENIUS THEORY

Acids – produce H+ ions (or hydronium ions H3O+) as the only positive ion.

HCl(l) Cl- + H+

A substance with a carboxyl group(COOH) looks like a base when you look at the chemical formula but it is an acid. (Acetic Acid = HC2H3O2 = CH3COOH)

CH3COOH + H2O CH3COO- + H+

ARRHENIUS THEORY

Bases – produce OH- ions (or hydroxide ions). Some bases DO NOT have hydroxide ions

attached. Amines – organic compounds containing C and N.

Amines are bases even though they do not have an hydroxide ion. Instead they react with water to produce the OH- ion.

NH3 + H2O NH4+ + OH-

~Caution~ Alcohols – contain an –OH but ARE NOT bases.

Example: CH3OH (hydroxyl group on a carbon chain)

TYPES OF ACIDS AND BASES

Acids Monoprotic: produce

one H+ ion. HCl

Diprotic: produce two H+ ions. H2SO4

Triprotic: produce three H+ ions. H3PO4

Bases Monohydroxy:

produce one OH- ion. NaOH

Dihydroxy: produce two OH- ions. Ba(OH)2

Trihydroxy: produce three OH- ions. Al(OH)3

STRENGTH OF ACIDS AND BASES

Determined by the amount of Ionization. Strong Acids

100% dissociation in water. Great conductors of electricity.

HNO3(aq) + H2O(l) H3O+(aq) + NO3

-(aq)

HI, HCl, HBr, H2SO4, and HClO4 are strong acids.

Weak Acids Much less than 100% dissociation. Poor conductors of electricity.

CH3COOH(aq) + H2O(l) CH3COO-(aq) + H3O+

(aq)

Acetic Acid(CH3COOH)

STRENGTH OF ACIDS AND BASES

Determined by the amount of Ionization. Strong Bases

100% dissociation (ionization) in water. Great conductors of electricity.

NaOH(aq) Na+(aq) + OH-

(aq)

KOH, Ca(OH)2, Group 1 or 2 metals with hydroxide!!

Weak Bases Much less than 100% dissociation (ionization). Poor conductors of electricity.

NH3(aq) + H2O(l) NH4+

(aq) + OH-(aq)

Ammonia (NH3)

BRØNSTED-LOWRY ACIDS AND BASES

Acids – Proton Donors According to the Brønsted-Lowy concept, an acid

is the chemical species that donates the proton in a proton transfer reaction.

Bases – Proton Acceptors According to the Brønsted-Lowy concept, a base

is the chemical species that accepts the proton in a proton transfer reaction.

A “proton” is really just a hydrogen that has lost its electron…H+

CONJUGATE PAIRS

The Brønsted-Lowry concept defines a species as an acid or a base according to its function in the proton-transfer reaction.

Consider the Reaction of NH3 + H2O

In the forward reaction, NH3 accepts a proton donated by H2O. Thus, NH3 is a base and H2O is an acid.

CONJUGATE PAIRS

In the reverse reaction, NH4+ donates a

proton to OH- which accepts it. Thus, NH4+ is

acid and OH- is the base.

CONJUGATE PAIRS

The species NH4+ and NH3 are a conjugate acid-

base pair. A conjugate acid-base pair consists of two species in an

acid-base reaction, one acid and one base, that differ by the loss or gain of a proton.

NH4+ is the conjugate acid of NH3

NH3 is the conjugate base of NH4+

The species OH- and H2O are a conjugate-acid base pair as well. OH- is the conjugate base of H2O H2O is the conjugate acid of OH-

CONJUGATE PAIRS

CONJUGATE PAIRS…PRACTICE

CONJUGATE PAIRS…PRACTICE PROBLEMS

Label the Acid, Base, Conjugate Acid, Conjugate Base in each reaction.

STRENGTH OF ACID-BASE CONJUGATE PAIRS

Strong Acids (Proton Donors) have weak conjugate bases.

Strong Bases (Proton Acceptors) have weak conjugate acids.

Strong acids and strong bases are always on the same side of an equation.

An acid can donate it H+ to any base EXCEPT it’s conjugate base.

Example: H3PO4 can donate to F-, but not to PO43-

PRACTICE PROBLEMS

Write the conjugate base for each.

1. HCl ______________

2. H2CrO4 ___________

3. NH4+ _____________

4. NH3 ______________

Write the conjugate acid for each.

1. F- _________________

2. H2PO4- ___________

3. NH3 ______________

4. HSO4- ____________

PRACTICE PROBLEMS

Write the conjugate base for each.

1. HCl _____Cl-_______

2. H2CrO4 __CrO42-__

3. NH4+ _____NH3_____

4. NH3 ____NH2-____

Write the conjugate acid for each.

1. F- ____HF____

2. H2PO4- _H3PO4_

3. NH3 ____NH4+ ____

4. HSO4- ___H2SO4___

PRACTICE PROBLEMS

1. CH3COO- + H30+ CH3COOH + H2O

2. HCl + H2O H3O+ + Cl-

3. NH2- + H2O NH3 + OH-

4. H3O + OH- H2O +H2O

5. CN- + H2O HCN + OH-

6. HClO4 + CH3COOH ClO4- + CH3COOH2

+

7. HCN + H2O H3O+ + CN-

PRACTICE PROBLEMS

1. CH3COO- + H30+ CH3COOH + H2O

2. HCl + H2O H3O+ + Cl-

3. NH2- + H2O NH3 + OH-

4. H3O + OH- H2O +H2O

5. CN- + H2O HCN + OH-

6. HClO4 + CH3COOH ClO4- + CH3COOH2

+

7. HCN + H2O H3O+ + CN-

ConjugateBase

ConjugateAcidAcidBase

BaseAcidConjugate

AcidConjugate

Base

AcidBaseConjugate

AcidConjugate

Base

BaseAcidConjugate

AcidConjugate

Base

AcidBaseConjugate

AcidConjugate

Base

BaseAcidConjugate

AcidConjugate

Base

Acid BaseConjugate

Acid Conjugate

Base

PRACTICE PROBLEMS

1. HSO4- + HCl H2SO4 +Cl-

2. SO42- + HNO3 HSO4

- + NO3-

3. NH4+ + HSO4

- NH3 + H2SO4

4. HCl + Al(H2O)5(OH)2+ Cl- + Al(H2O)63+

5. NH3 + NH3 NH4+ + NH2

-

PRACTICE PROBLEMS

1. HSO4- + HCl H2SO4 +Cl-

2. SO42- + HNO3 HSO4

- + NO3-

3. NH4+ + HSO4

- NH3 + H2SO4

4. HCl + Al(H2O)5(OH)2+ Cl- + Al(H2O)63+

5. NH3 + NH3 NH4+ + NH2

-

StrongAcid

StrongBase

WeakAcid

WeakBase

WeakBase

StrongAcid

StrongBase

WeakAcid

WeakBase

WeakAcid

StrongBase

StrongAcid

StrongAcid

StrongBase

WeakBase

WeakAcid

WeakBase

WeakAcid

StrongAcid

StrongBase

AMPHOTERIC (AMPHIPROTIC) SPECIES

Substances that act as both an acid or a base.

Depends on chemical environment. Examples: H2O, HSO4

-, HS-

In the reaction between NH3 and H2O, water is an acid. In the reaction between HNO2 and H2O, water is a

base. Water (H2O) is an amphoteric substance.

LEWIS ACIDS AND BASES

Lewis Acid – a substance that ACCEPTS an electron(e-) pair.

Lewis Base – a substance that DONATES an electron(e-) pair.

Formation of the Hydronium Ion is an excellent example.

LEWIS ACID/BASE REACTION

REACTIONS INVOLVING ACIDS

Steps…1) Check the metal on Table J. If it is above H2 proceed.

2) Write H2 as a product.

3) Combine the metal with the negative (-) ion to form an ionic salt. (write the metal first, followed by the negative ion.)

4) Assign charges5) Criss Cross, if necessary6) Balance the equation.

REACTIONS INVOLVING ACIDS, EXAMPLES…

1. HCl + Sr __________ + __________

2. H3PO4 + Zn __________ + __________

3. HNO3 + Au __________ + __________

4. HC2H3O2 + K __________ + __________

5. HF + Cu __________ + __________


1. HCl + Sr __________ + __________ Steps…

1) Check the metal on Table J. If it is above H2 proceed.


1. HCl + Sr __________ + _____H2(g)_____ Steps…




1. HCl + Sr __________ + _____H2(g)_____ Steps…




4) Assign charges


1. HCl + Sr ___SrCl2____ + _____H2(g)_____ Steps…




4) Assign charges5) Criss Cross, if necessary


1. HCl + Sr ___SrCl2____ + _____H2(g)_____ Steps…




4) Assign charges5) Criss Cross, if necessary6) Balance the equation.


1. 2HCl + Sr ___SrCl2____ + _____H2(g)_____

2. H3PO4 + Zn __________ + __________

3. HNO3 + Au __________ + __________

4. HC2H3O2 + K __________ + __________

5. HF + Cu __________ + __________


1. 2HCl + Sr ___SrCl2____ + _____H2(g)_____

2. 2H3PO4 + 3Zn _Zn3(PO4)2_ + ___3H2(g)___

3. HNO3 + Au __AuNO3__ + _____H2(g)_____

4. HC2H3O2 + K _KC2H3O2_ + _____H2(g)_____

5. HF + Cu _____CuF_____ + _____H2(g)_____

NEUTRALIZATION

Acid + Base Salt + H2O

NEUTRALIZATION

Acid + Base Salt + H2O Steps…

1) Form water, H2O.

2) Get rid of all H+ on the acid, and all OH- on the base.

3) Write the metal 1st and the negative(-) ion 2nd.

4) Assign Charges.

5) CrissCross, if necessary.

6) Balance the equation.

NEUTRALIZATION…EXAMPLES

1. CH3COOH + NaOH __________ + __________

2. KOH + HCl __________ + __________

3. HCl + NaOH __________ + __________

4. H2SO4 + NaOH __________ + __________

5. HCl + Ba(OH)2 __________ + __________

6. HNO3 + LiOH __________ + __________

NEUTRALIZATION

CH3COOH + NaOH __________ + __H2O__ Steps…

1) Form water, H2O.



4) Assign Charges.



NEUTRALIZATION


1) Form water, H2O.



4) Assign Charges.



CH3COO + Na

NEUTRALIZATION


1) Form water, H2O.



4) Assign Charges.



Na CH3COO

CH3COO + Na

NEUTRALIZATION


1) Form water, H2O.



4) Assign Charges.



Na CH3COO

CH3COO + Na

Na+ CH3COO-

NEUTRALIZATION

CH3COOH + NaOH NaCH3COO + __H2O__ Steps…

1) Form water, H2O.



4) Assign Charges.



Na CH3COO

CH3COO + Na

Na+ CH3COO-

NaCH3COO

NEUTRALIZATION…EXAMPLES

1. CH3COOH + NaOH NaCH3COO + __H2O__

2. KOH + HCl __________ + __________

3. HCl + NaOH __________ + __________

4. H2SO4 + NaOH __________ + __________

5. HCl + Ba(OH)2 __________ + __________

6. HNO3 + LiOH __________ + __________

SPECTATOR IONS

Ions found on both sides of the equation that are not involved in making water. Not part of the Net Arrhenius Equation.

Spectator Ions for the previous example…

CH3COO-, Na+

NET ARRHENIUS EQUATION

Does NOT include spectator ions!!

Net Arrhenius Equation is always…

H+ + OH- H2O

THE PH SCALE

The pH Scale expresses the strength of acids and bases. Logarithmic Scale – one jump on

the scale represents a tenfold change in [H+]

[ ] = concentration (usually Molarity)

pH > 7 is a Base pH = 7 is Neutral pH < 7 is an Acid

As pH ↑ [H+] ↓

The stronger the acid the more H+ ions it produces.

The stronger the base the more OH- ions it produces.

PH EXAMPLES A solution with a pH of 1 has how many times the

amount of H+ compared to a solution with a ph of 6?

A solution with a pH of 2 has how many times the amount of H+ compared to a solution with a pH of 5?

If the [H+] increases, the [OH-] decreases by the same amount. If the pH changes from 8 to 13, the [H+] decreases _______

times and the [OH-] increases _______ times. If the pH changes from 6 to 2, the [H+] increases _______

times and the [OH-] decreases _______ times.






One jump on the scale represents a tenfold change in [H+].



6 – 1 = 5







6 – 1 = 5





Moving from one to six is equivalent to five jumps on the pH scale.



6 – 1 = 5






Use the difference as an exponent to

ten.



6 – 1 = 5






Use the difference as an exponent to

ten.

105 = 100,000 times more H+



6 – 1 = 5

105 = 100,000 times more H+







6 – 1 = 5

105 = 100,000 times more H+


5 – 2 = 3

103 = 1,000 times more H+






6 – 1 = 5

105 = 100,000 times more H+


5 – 2 = 3

103 = 1,000 times more H+

If the [H+] increases, the [OH-] decreases by the same amount. If the pH changes from 8 to 13, the [H+] decreases 100,000

times and the [OH-] increases 100,000 times. If the pH changes from 6 to 2, the [H+] increases 10,000

times and the [OH-] decreases 10,000 times.

CALCULATING THE PH OF A SUBSTANCE

pH = -log [H+]Recall that the [ ] mean concentration (usually Molarity)

Example: If [H+] = 1.0 x 10-10 what is the pH?pH = -log [H+]pH = -log (1.0 x 10-10)pH = 10

Example: If [H+] = 1.8 x 10-5 what is the pH?pH = -log [H+]pH = -log (1.0 x 10-5)pH = 4.74


Find the pH of these:

1) A 0.15M solution of HCl.

2) A 3.00 x 10-7M solution of HNO3




A 0.15M solution of HCl has the same concentration for each of its respective ions.





0.15M H+

0.15M Cl-




pH = -log [H+]pH = -log (0.15)pH = pH = 0.82


0.15M H+

0.15M Cl-

[H+] = 0.15M



1) A 0.15M solution of HCl.pH = -log [H+]pH = -log (0.15)pH = 0.82






A 3.00 x 10-7M solution of HNO3 has the same concentration for each of its respective ions.






0.15M H+

0.15M NO3-






0.15M H+

0.15M NO3-

[H+] = 3.00 x 10-7





pH = -log [H+]pH = -log (3.00 x 10-7)pH = 6.5

PH CALCULATIONS…SOLVING FOR H+

If the pH of Diet Coke is 3.12, what is the [H+]? Since pH = -log [H+] then…

[H+] = 10-pH

[H+] = 10-pH

[H+] = 10-3.12

[H+] = 7.6 x 10-4M


A solution has a pH of 8.5, what is the molarity of hydrogen ions in the solution?



[H+] = 10-pH



[H+] = 10-pH

[H+] = 10-8.5



[H+] = 10-pH

[H+] = 10-8.5

[H+] = 3.2 x 10-9 M

MORE ABOUT WATER

H2O can function as both an ACID and a BASE. Amphoteric substance

In pure water there can be Autoionization.

Equilibrium Constant for water = Kw

Kw = [H3O+][OH-] = 1.00 x 10-14 at 25oC In a neutral solution [H3O+] = [OH-] So… [H3O+] = [OH-] = 1.00 x 10-7

http://en.wikipedia.org/wiki/File:Autoprotolyse_eau.svg

POH

Strong Acids and Strong Bases are opposites. pH and pOH are opposites as well.

pOH scale does not really exist, but it is useful for determining the pH of a base.

pOH = -log [OH-]

Since pH and pOH are on opposite ends,

pH + pOH = 14

[H30+], [OH-], PH AND POH

What is the pH of a 0.0010M NaOH solution?



A 0.0010M solution of NaOH has the same concentration for each of its respective ions.




0.0010M Na+

0.0010M OH-




0.0010M Na+

0.0010M OH-

[OH-] = 0.0010M




0.0010M Na+

0.0010M OH-

[OH-] = 0.0010M

[OH-] = 0.0010 (or 1.0 x 10-

3)pOH = -log [OH-]pOH = -log (0.0010)pOH = 3pH + pOH =14pH + 3 = 14pH = 11

~or~




0.0010M Na+

0.0010M OH-

[OH-] = 0.0010M

[OH-] = 0.0010 (or 1.0 x 10-

3)pOH = -log [OH-]pOH = -log (0.0010)pOH = 3pH + pOH =14pH + 3 = 14pH = 11

Kw = [H30+][OH-]1.00 x10-14 = [H30+](1.00 x10-3)[H30+] = 1.00 x10-11MpH = -log [H30+]pH = -log (1.00 x10-11)pH = 11`

~or~

[H30+], [OH-], PH AND POH…PROBLEMS The pH of rainwater collected in a certain region of the

northeastern United States on a particular day was 4.82; What is the H+ ion concentration of the rainwater?

The OH- ion concentration of a blood sample is 2.5 x10-7 M. What is the pH of the blood?



pH = -log [H+] so…

[H+] = 10-pH

[H+] = 10-4.82

[H+] = 1.5 x10-5 The OH- ion concentration of a blood sample is 2.5

x10-7 M. What is the pH of the blood?



pH = -log [H+] so…

[H+] = 10-pH

[H+] = 10-4.82

[H+] = 1.5 x10-5 The OH- ion concentration of a blood sample is 2.5 x10-7 M.

What is the pH of the blood?

pOH = -log [OH-]

pOH = -log (2.5 x10-7)

pOH = 6.6

pOH + pH = 14

6.6 + pH = 14

pH = 7.4

GENERAL SUMMARY OF FORMULAS

pH = -log [H+]

pOH = -log [OH-]

[H+][OH-] = 1.0 x10-14

[H+] = 10-pH

[OH-] = 10-pH

pH + pOH = 14

CALCULATING [H30+], [OH-], PH AND POH

Problem 1: A chemist dilutes concentrated hydrochloric acid to make two solutions: (a) 3.0M and (b) 0.0024M. Calculate the [H3O+], pH, [OH-] and pOH of the two solutions at 25oC.


Problem 1: A chemist dilutes concentrated hydrochloric acid to make two solutions: (a) 3.0M and (b) 0.0024M. Calculate the [H3O+], pH, [OH-] and pOH of the two solutions at 25oC.

Solution (a)

[H3O+] = 3.0MpH = -log [H3O+] pH = -log (3.0)pH = -0.48pH + pOH = 14-0.48 + pOH = 14pOH = 14.48[OH-] = 10-pOH

[OH-] = 10-14.48

[OH-] = 3.31 x10-15

Solution (b)

[H3O+] = 0.0024MpH = -log [H3O+] pH = -log (0.0024)pH = 2.62pH + pOH = 142.62 + pOH = 11.38pOH = 11.38[OH-] = 10-pOH

[OH-] = 10-11.38

[OH-] = 4.17 x10-12


Problem 2: What is the [H30+], [OH-] and pOH of a solution with pH = 3.67? Is this an acid, base, or neutral?



pH = 3.67 = acid[H3O+] = 10-pH

[H3O+] = 10-3.67

[H3O+] = 2.14 x10-4

pH + pOH = 143.67 + pOH = 14pOH = 10.33[OH-] = 10-pOH

[OH-] = 10-10.33

[OH-] = 4.68 x10-11



pH = 8.05 = base[H3O+] = 10-pH

[H3O+] = 10-8.05

[H3O+] = 8.91 x10-9MpH + pOH = 148.05 + pOH = 14pOH = 5.95[OH-] = 10-pOH

[OH-] = 10-5.95

[OH-] = 1.12 x10-6M

USING TABLE M

USING TABLE M

Used to approximate the pH of a substance.

Indicators tend to have a distinct color at the ends of their useful pH range. So a solution with methyl orange indicator is red

at a pH of 0.

USING TABLE M



at a pH of 0. As a base is added pH begins to rise.

USING TABLE M



at a pH of 0. As a base is added pH begins to rise. At a pH of about 3.2 the red color begins to

change yellow but you first enter a intermediate color region that is a mixture of these two colors.

USING TABLE M



at a pH of 0. As a base is added pH begins to rise. At a pH of about 3.2 the red color begins to

change yellow but you first enter a intermediate color region that is a mixture of these two colors.

At a pH of about 4.4 the intermediate color region ends and the solution is now completely yellow all the way through to pH of 14.

USING TABLE M…QUESTIONS What color will bromcresol green

be in an acid with a pH of 2.2?

What color will bromcresol green be in a base with a pH of 10.0?

Can bromcresol green determine whether an unknown substance is an acid or base? Why or why not?

What is the pH range for a substance that turns methyl orange yellow and bromothymol blue yellow?

USING TABLE M…QUESTIONS What color will bromcresol green

be in an acid with a pH of 2.2?Yellow


Can bromcresol green determine whether an unknown substance is an acid or base? Why or why not?

What is the pH range for a substance that turns methyl orange yellow and bromothymol blue yellow?

USING TABLE M…QUESTIONS What color will bromcresol green be

in an acid with a pH of 2.2?Yellow


Blue Can bromcresol green determine

whether an unknown substance is an acid or base? Why or why not? No, because the solution would be totally blue at a pH of 5.4 which is still in the acid range and would remain blue throughout the basic range

What is the pH range for a substance that turns methyl orange yellow and bromothymol blue yellow? 4.4 - 6.0

TITRATION

TITRATION

Neutralization Reaction

H2C2O4(aq) + 2NaOH(aq) Na2C2O4(aq) + 2H2O(l)

Carry out this reaction using a TITRATION.

In neutralization reactions the # of moles of H+

must equal the # of moles of OH-.

Acid Base Salt Water

SET UP FOR TITRATING AN ACID WITH A BASE

Flask containing aqueous solution of sample being analyzed.

50-mL buret containing aqueous NaOH of accurately known concentration.

The NaOH in the buret is slowly added to the sample being analyzed.

SET UP FOR TITRATING AN ACID WITH A BASE

50-mL buret containing aqueous NaOH of accurately known concentration.

The NaOH in the buret is slowly added to the sample being analyzed.

When the amount of NaOH added from the buret exactly equals the amount of H+ supplied by the acid being analyzed, the dye (indicator) changes color.

Flask containing aqueous solution of sample being analyzed.

TITRATION

1. Pour standard solution into the buret. (You know the M of this solution.)

2. Add the standard solution to the solution of unknown M in the Erlenmeyer flask below.

3. Indicator changes color when the end point is reached. [H+] = [OH-] at the equivalence point.

TITRATION FORMULA

#MAVA = #MBVB

# = the number of H+ or OH- produced per mole of acid or base.

M = MolarityV = VolumeSubscriptsA or B = Acid or Base respectively.

Example: For H2CO3 you would plug in 2 for the # on the acid side.

TITRATION…PRACTICE PROBLEM 1 35.6 mL of NaOH is neutralized with 25.2 mL of 0.0998

M HCl by titration to an equivalence point. What is the concentration of the NaOH?



#MAVA = #MBVB



#MAVA = #MBVB

(1)(0.0998 M)(25.2 mL) = (1)MB(35.6 mL)



#MAVA = #MBVB

(1)(0.0998 M)(25.2 mL) = (1)MB(35.6 mL)

MB = 0.071 M

TITRATION…PRACTICE PROBLEM 1

60.3 mL of Ca(OH)2 is neutralized with 33.2 mL of 1.0 M HF by titration to an equivalence point. What is the concentration of the Ca(OH)2?



#MAVA = #MBVB



#MAVA = #MBVB

(1)(1.0 M)(33.2 mL) = (2)MB(60.3 mL)



#MAVA = #MBVB

(1)(1.0 M)(33.2 mL) = (2)MB(60.3 mL)

MB = 0.275 M

Acids and Bases

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