chapter #14 acids, bases, and salts. acids, bases and salts topics the arrhenius theory the...

Chapter #14Acids, Bases, and Salts

Acids, Bases and Salts TopicsThe Arrhenius Theory The Brønsted Theory Naming Acids (See Nomenclature Notes) The Self-Ionization of Water The pH Concept Properties of Acids Properties of Bases Salts The Strengths of Acids and BasesAnalyzing Acids and BasesTitration CalculationsHydrolysis Reactions of SaltsBuffers

History of Acids and BasesIn the early days of chemistry chemists were organizing physical and chemical properties of substances. They discovered that many substances could be placed in two different property categories:

Substance A

1. Sour taste

2. Reacts with carbonates to make CO2

3. Reacts with metals to produce H2

4. Turns blue litmus pink

5. Reacts with B substances to make salt water

Substance B

1. Bitter taste

2. Reacts with fats to make soaps

3. Do not react with metals

4. Turns red litmus blue

5. Reacts with A substances make salt and water

Arrhenius was the first person to suggest a reason why substances are in A or B due to their ionization in water.

The Swedish chemist Svante Arrhenius proposed the first definition of acids and bases.(Substances A and B becameknown as acids and bases)According to the Arrhenius model:

“acids are substances that dissociate in water to produce H+ ions and bases are substances that dissociate in water to produce OH- ions”

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

HCl (aq) H+ (aq) + Cl- (aq) Acid

Arrhenius Theory

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What is H+?

+e-

+

Hydrogen (H) Proton (H+)

Unknown to Arrhenius free H+ ions do not exist in water. They covalently react with water to produce hydronium ions, H3O+.

or:H+ (aq) + H2O (l) H3O+ (aq)

This new bond is called a coordinate covalent bond since both new bonding electrons come from the same atom

Hydronium Ion

Hydronium ion is the name for H3O+ and is often times abbreviated as H+ (aq) they both mean the same thing.

What is the difference between a strong acid and a weak acid?

Hydronium Ion


What is the difference between a strong acid and a weak acid? Strong acids ionize 100% and weak ones do not!

Hydronium Ion



A single arrow is used to represent the ionization of a strong acid. Double arrows (Equilibrium) are used to represent weak acids.

For example: HCl (g) H+ (aq) + Cl - (aq)

HF (g) H+ (aq) + F -

Hydronium Ion





HF (g) H+ (aq) + F -

According to Arrhenius, is water an acid or base?HOH (l) H+ (aq) + OH – (aq)

Hydronium Ion





HF (g) H+ (aq) + F -

According to Arrhenius, is water an acid or base?HOH (l) H+ (aq) + OH – (aq)

Neither, he called it Neutral (same amount of OH- and H+

Hydronium Ion

Strong Acids and Bases

How can we identify strong acids or bases?


How can we identify strong acids or bases?Easy memorize them!

How can we identify strong acids or bases?Easy, memorize them!

Memorized Strong Acids1. HClO4

2. H2SO4

3. HI

4. HBr

5. HCl

6. HNO3

Memorized Strong BasesHydroxides of group 1 and 2metals, excluding Be and Mg


Johannes Brønsted and Thomas Lowry revised Arrhenius’s acid-base theory to include this behavior. They defined acids and bases as follows:

“An acid is a hydrogen containing species that donates a proton. A base is any substance that accepts a proton”

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

In the above example what is the Brønsted acid? What is the Brønsted base?

Bronsted Lowry

Bronsted Lowry Theory

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http://www.chemcool.com/biography/Bronsted.gif

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

In reality, the reaction of HCl with H2O is an equilibrium and occurs in both directions, although in this case the equilibrium lies far to the right.

For the reverse reaction Cl - behaves as a Brønsted base and H3O+ behaves as a Brønsted acid.

The Cl- is called the conjugate base of HCl. Brønsted acids and bases always exist as conjugate acid-base pairs.

Bronsted Lowry Theory

In pure water (no solute) water molecules behave as both an acid and base!!

e.g.

H2O (l) + H2O (l) H3O+ (aq) + OH- (aq)

This is called the self-ionization (autoionizaion) of water. Although the equilibrium lies far to the left it is very important to take into consideration, especially for living systems.

Does anyone know how we write the equilibrium constant for this reaction?

Autoionization of Water

The auto-ionization of water is described by the equation:

H2O (l) + H2O (l) H3O+ (aq) + OH- (aq)The equilibrium constant for this reaction is given by:

]OH][O3H[2]O2H[K

2]O2H[

]OH][O3H[

]O2H][O2H[

]OH][O3H[K

Kw = K[H2O]2 = 10-14 This equilibrium lies very much to the left i.e. mostly water. For pure water [OH-] = [H+] = 1 x 10-7 M


]OH][O3H[wK2]O2H[K

2]O2H[

]OH][O3H[

]O2H][O2H[

]OH][O3H[K

As [OH-] and [H+] are so small the [H2O] is not affected by their formation. It is useful to define a new constant Kw such that:

Kw is called the ion product of water.

What is the value for the ion product of water?

1.00 gml 18.0 g

mole ml

10-3 L= 55.5 M


]OH][O3H[wK2]O2H[K

2]O2H[

]OH][O3H[

]O2H][O2H[

]OH][O3H[K

As [OH-] and [H+] are so small the [H2O] is not affected by their

formation. It is useful to define a new constant Kw such that:

Kw is called the ion product of water.

What is the value for the ion product of water?

[H+][OH-] = 10-14

1.00 gml 18.0 g

mole ml

10-3 L= 55.5 M


We define an aqueous solution as being neutral when the

[H+] = [OH-].

We define an aqueous solution as being acidic when

[H+] > [OH-].

We define an aqueous solution as being basic when

[H+] < [OH-].

However, in each case Kw = 1 x 10-14 M2

[H+] = 0.0000001 = 10-7 (how can this be abbreviated further?)



[H+] = [OH-].


[H+] > [OH-].


[H+] < [OH-].



By just describing the power



[H+] = [OH-].


[H+] > [OH-].


[H+] < [OH-].



By just describing the power Called the power of H, or pH.



[H+] = [OH-].


[H+] > [OH-].


[H+] < [OH-].



By just describing the power Called the power of H, or pH.

pH = 7Our math departments tells us that log means power too.


The mathematical definition of pH using [H+] for [H3O+] is listed below:

pH = -log [H+], or [H+]= 1x10-pH (both are mathematically

equivalent)

How about the power for the OH -, what should this be called?


pH = -log [H+], or [H+] = 1x10-pH (both are mathematically equivalent)

How about the power for the OH -, what should this be called? Would you believe pOH?



pH = -log [H+], or [H+]= 1x10-pH (both are mathematically equivalent)


Have you heard of pOH before?



pH = -log [H+], or [H+]= 1x10-pH (both are mathematically equivalent)


Have you heard of pOH before?

pH + pOH = 14 for water solutions.


Now for some examples1. Find the pH and pOH, when [H+] = 10-4

Now for some examples1. Find the pH and pOH, when [H+] = 10-4

pH = 4 and pOH = 10, since they must add to 14

using the calculator pH = -log [H+], type in 10-4, push the log button and pH = -(-4) = 4. Same for pOH

A pH Number lineNumber lines have been used in history and math classes, so to keep up we use them in chemistry classes.

pH = 16

pH = 12

pH = 7

pH = 2[H+] = 10-2


pH = 16

pH = 12

pH = 7

pH = 2[H+] = 10-2

[OH -] = 10-12


pH = 16

pH = 12

pH = 7

pH = 2[H+] = 10-2

[OH -] = 10-12

[H+] > [OH -]acidic


pH = 16

pH = 12

pH = 7

pH = 2[H+] = 10-2

[OH -] = 10-12

[H+] = 10-7

[OH -] = 10-7

[H+] > [OH -]acidic


pH = 16

pH = 12

pH = 7

pH = 2[H+] = 10-2

[OH -] = 10-12

[H+] = 10-7

[OH -] = 10-7

[H+] = [OH -]neutral

[H+] > [OH -]acidic


pH = 16

pH = 12

pH = 7

pH = 2[H+] = 10-2

[OH -] = 10-12

[H+] = 10-7

[OH -] = 10-7

[H+] =10-12

[OH -] = 10-2


[H+] > [OH -]acidic


pH = 16

pH = 12

pH = 7

pH = 2[H+] = 10-2

[OH -] = 10-12

[H+] = 10-7

[OH -] = 10-7

[H+] =10-12

[OH -] = 10-2

[H+] < [OH -]basic


[H+] > [OH -]acidic


pH = 16

pH = 12

pH = 7

pH = 2[H+] = 10-2

[OH -] = 10-12

[H+] = 10-7

[OH -] = 10-7

[H+] =10-12

[OH -] = 10-2

[H+] =10-16

[OH -] =

[H+] < [OH -]basic


[H+] > [OH -]acidic


pH = 16

pH = 12

pH = 7

pH = 2[H+] = 10-2

[OH -] = 10-12

[H+] = 10-7

[OH -] = 10-7

[H+] =10-12

[OH -] = 10-2

[H+] =10-16

[OH -] = 102

[H+] < [OH -]basic


[H+] > [OH -]acidic


pH = 16

pH = 12

pH = 7

pH = 2[H+] = 10-2

[OH -] = 10-12

[H+] = 10-7

[OH -] = 10-7

[H+] =10-12

[OH -] = 10-2

[H+] =10-1

[OH -] = 102

[H+] < [OH -]basic

[H+] < [OH -]basic


[H+] > [OH -]acidic


pH = 16

pH = 12

pH = 7

pH = 2[H+] = 10-2

[OH -] = 10-12

[H+] = 10-7

[OH -] = 10-7

[H+] =10-12

[OH -] = 10-2

[H+] =10-16

[OH -] = 102

[H+] < [OH -]basic

[H+] < [OH -]basic


[H+] > [OH -]acidic

acidic


pH = 16

pH = 12

pH = 7

pH = 2[H+] = 10-2

[OH -] = 10-12

[H+] = 10-7

[OH -] = 10-7

[H+] =10-12

[OH -] = 10-2

[H+] =10-16

[OH -] = 102

[H+] < [OH -]basic

[H+] < [OH -]basic


[H+] > [OH -]acidic

acidic

basic

Acids undergo characteristic double replacement reactions

with oxides, hydroxides, carbonates and bicarbonates.

e.g.

2HCl (aq) + CuO (s) CuCl2 (aq) + H2O (l)

2HCl (aq) + Ca(OH)2 (aq) CaCl2 (aq) + 2H2O (l)

2HCl (aq) + CaCO3 (aq) CaCl2 (aq) + H2O (l) + CO2 (g)

2HC l (aq) + Sr(HCO3)2 (aq) SrCl2 (aq) + 2H2O (l) + 2CO2 (g)

Equations With Acuids

Bases undergo a double replacement reaction with acids

called neutralization:

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

In words this well known reaction is often described as:

“acid plus base = salt plus water”

We previously discussed this reaction when describing types of

reactions.

Equations With Acuids

We have discussed the double replacement reactions and ionic

equations before. Since the acids and bases undergo double

replacement reactions called neutralization reactions, then they

can have ionic equations too.

e.g.

Formula equation:

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

Ionic equation:

H+ (aq) + Cl- (aq) + Na+ (aq) + OH- (aq) Na+ (aq) + Cl- (aq) + H2O (l)

Net ionic equation:

H+ (aq) + OH- (aq) H2O (l)

Ionic Equations (a review)

Another property of acids is their reaction with certain metals to

produce hydrogen gas, H2 (g).

Zn (s) + 2HC l (aq) H2 (g) + ZnCl2 (aq)

This is an example of a single replacement reaction and is a

redox reaction.

Total ionic equation:

Zn (s) + 2H+ (aq) + 2Cl- (aq) H2 (g) + Zn2+ (aq) + 2Cl- (aq)

Net ionic equation:

Zn (s) + 2H+ (aq) H2 (g) + Zn2+ (aq)

Acidic Single Replacement Reactions

SaltsSalts are the ionic product of an acid base neutralization reaction.

Acidic Salts are formed from a strong acid and a weak base.

Neutral salts are formed from a strong acid and strong base.

Basic salts are formed from a strong base and a weak acid.

Give the acid and base the following salts were formed from and label the salts as acidic, basic, or neutral.

1. NaCl

2. NaC2H3O2

3. NH4Cl






1. NaCl

1. NaC2H3O2

1. NH4Cl

NaCl + HOHReactants are?






1. NaCl

2. NaC2H3O2

3. NH4Cl

NaCl + HOHHCl + NaOH

NaC2H3O2 + HOH

S.A. s.b.






1. NaCl

2. NaC2H3O2

3. NH4Cl


NaC2H3O2 + HOH

S.A. s.b.Neutral salt






1. NaCl

2. NaC2H3O2

3. NH4Cl


NaC2H3O2 + HOHHC2H3O2 + NaOH

Neutral salt s.a. s.b.






1. NaCl

2. NaC2H3O2

3. NH4Cl


NaC2H3O2 + HOHHC2H3O2 + NaOHw.a. s.b.







1. NaCl

2. NaC2H3O2

3. NH4Cl


NaC2H3O2 + HOHHC2H3O2 + NaOHw.a. s.b.basic salt







1. NaCl

2. NaC2H3O2

3. NH4Cl




NH4Cl + HOH






1. NaCl

2. NaC2H3O2

3. NH4Cl




NH4Cl + HOHNH4OHHCl +






1. NaCl

2. NaC2H3O2

3. NH4Cl




NH4Cl + HOHNH4OHHCl +s.a. w.b.






1. NaCl

2. NaC2H3O2

3. NH4Cl



neutral salt s.a. s.b.

NH4Cl + HOHNH4OHHCl +s.a. w.b.acidic salt

Acid, Base, and Salt Hydrolysis

HBr (aq)


HBr (aq) H+ (aq) + Br - (aq)


HBr (aq) H+ (aq) + Br - (aq) Acidic, because H+ (aq)



0.1 Initial concentration



0.1 Initial concentration0.0



0.1 Initial concentration0.0 ?



0.1 Initial concentration0.0 0.0




Final concentration?




Final concentration0.0




Final concentration0.0 ?




Final concentration0.0 0.1




Final concentration0.0 0.1 ?




Final concentration0.0 0.1 0.1





[H+] = ?





[H+] = 0.1





pH = ?

[H+] = 0.1





pH = ?

[H+] = 0.1 = 10-1





pH = 1

[H+] = 0.1 = 10-1





pH = 1

Ca(OH)2 (aq)





pH = 1

Ca(OH)2 (aq) Ca2+ (aq) + 2 OH- (aq)





pH = 1

Ca(OH)2 (aq) Ca2+ (aq) + 2 OH- (aq) acidic?





pH = 1

Ca(OH)2 (aq) Ca2+ (aq) + 2 OH- (aq) No, basic OH-





pH = 1

Ca(OH)2 (aq) Ca2+ (aq) + 2 OH- (aq) Initial concentration0.1





pH = 1

Ca(OH)2 (aq) Ca2+ (aq) + 2 OH- (aq) Initial concentration0.1 0.0





pH = 1

Ca(OH)2 (aq) Ca2+ (aq) + 2 OH- (aq) Initial concentration0.1 0.0 ?





pH = 1

Ca(OH)2 (aq) Ca2+ (aq) + 2 OH- (aq) Initial concentration0.1 0.0 0.0





pH = 1


Final concentration?





pH = 1


Final concentration0.0 ?





pH = 1


Final concentration0.0 0.1





pH = 1


Final concentration0.0 0.1 ?





pH = 1







pH = 1



[OH - ] = ?





pH = 1



[OH - ] = 0.2





pH = 1



pOH = ?

[OH - ] = 0.2





pH = 1



pOH = - log[OH-]

[OH - ] = 0.2





pH = 1



pOH = - log[OH-] = - log[0.2]

[OH - ] = 0.2





pH = 1



pOH = - log[OH-] = - log[0.2] = -(-0.698970004)

pOH = 0.7

[OH - ] = 0.2





pH = 1



pOH = - log[OH-] = - log[0.2] = -(-0.698970004)

pOH = 0.7

pH = ?

[OH - ] = 0.2





pH = 1



pOH = - log[OH-] = - log[0.2] = -(-0.698970004)

pOH = 0.7

pH = 14.0 - 0.07 = 13.3

[OH - ] = 0.2





pH = 1


final concentration0.0 0.1 0.2

pH = 13.3





pH = 1



pH = 13.3

NaF (aq)





pH = 1



pH = 13.3

NaF (aq) Na+ (aq) + F – (aq)





pH = 1



pH = 13.3

NaF (aq) Na+ (aq) + F – (aq) Acidic, basic, or neutral?





pH = 1



pH = 13.3

NaF (aq) Na+ (aq) + F – (aq)Basic, since HF is w.a. and NaOH is s.b.

Will sodium and fluorine ions react with water?





pH = 1



pH = 13.3



Na+ + HOH NaOH (sb) + H+





pH = 1



pH = 13.3



Na+ + HOH NaOH (sb) + H+

Na+ + HOH Na+ + OH- + H+

HOH OH- + H+ No Reaction, water cannot make water





pH = 1



pH = 13.3



Na+ + HOH NaOH + H+Cannot make strong acids or bases from weak oness.b.





pH = 1



pH = 13.3



Na+ + HOH NaOH + H+Cannot make strong acids or bases from weak ones

F - + HOH HF + OH-

w.a.





pH = 1



pH = 13.3



Na+ + HOH NaOH + H+Cannot make strong acids or bases from weak ones

F - + HOH HF + OH- Yes, HF weak acid and OH- is formed, thus basic salt!

w.a.


NH4Cl (aq) NH4+

(aq) + Cl- (aq)


NH4Cl (aq) NH4+

(aq) + Cl- (aq) acidic, basic, or neutral?


NH4Cl (aq) NH4+


HCl + NH4OH NH4Cl + HOH


NH4Cl (aq) NH4+


HCl + NH4OH NH4Cl + HOHs.a. w.b.


NH4Cl (aq) NH4+

(aq) + Cl- (aq) Acidic!


Will the ions from the salt combine with water?

NH4+ + HOH NH4OH + H+


NH4Cl (aq) NH4+





w.b.


NH4Cl (aq) NH4+





w.b.

This reaction is OK, since a w.b. is formed


NH4Cl (aq) NH4+





w.b.


Cl- + HOH HCl + OH-


NH4Cl (aq) NH4+





w.b.


Cl- + HOH HCl (sa) + OH-


NH4Cl (aq) NH4+





w.b.


Cl- + HOH H+ + Cl- + OH-


NH4Cl (aq) NH4+





w.b.


HOH H+ + OH-Again water cannot make water! NR


NH4Cl (aq) NH4+





w.b.


Cl- + HOH HCl + OH-

s.a.


NH4Cl (aq) NH4+





w.b.


Cl- + HOH HCl + OH-

s.a.

Cannot form s.a. from weaker reactants, thus N.R.


NH4Cl (aq) NH4+





w.b.


Cl- + HOH HCl + OH-

s.a.

Cannot form s.a. from weaker reactants, thus N.R.

Since H+ was formed in the first reaction, then [H+] is now

greater than [OH-] making the solution acidic


NaCl (aq)


NaCl (aq) Na+ (aq) + Cl- (aq)


NaCl (aq) Na+ (aq) + Cl- (aq) Acidic, basic, or neutral?



HCl + NaOH NaCl + HOH



HCl + NaOH NaCl + HOHs.a. s.b.


NaCl (aq) Na+ (aq) + Cl- (aq) Neutral!





Now react each of the ions with water.

Na+ + HOH

NaOH + H+





Na+ + HOH

NaOH + H+

s.b.





Na+ + HOH

NaOH + H+

s.b.

Cannot form strong bases from weaker ones, thus N.R.





Na+ + HOH

NaOH + H+

s.b.


Cl- + HOH HCl + OH-





Na+ + HOH

NaOH + H+

s.b.


Cl- + HOH HCl + OH-s.a.





Na+ + HOH

NaOH + H+

s.b.


Cl- + HOH HCl + OH-s.a.

Cannot form strong acids from weaker ones, thus N.R.

Buffers

Buffers are extremely important in chemistry and biology. They maintain a nearly consistent pH in various solutions.

BuffersBuffers are extremely important in chemistry and biology. They maintain a nearly consistent pH in various solutions.

Our blood must maintain a pH around 7.35-7.45. If the pH is above 7.45 you would have a condition called alkalosis. If the pH is below 7.35, then one would suffer from acidosis.



Acidosis leads to depression of the nervous system. Mild acidosis can result in dizziness, disorientation, or fainting; a more severe case can cause coma, or death.



What would happen to the pH of our blood if we were to eat acidic foods, such as apples, oranges, or limes? What might happen to the pH of our blood if some of the hydrochloric acid from our stomach were to seep into our blood?




What would happen to the pH of our blood if we were to eat acidic foods, such as apples, oranges, or limes? What might happen to the pH of our blood if some of the hydrochloric acid from our stomach were to seep into our blood? The pH would be lower in both


Despite the possibility of pH increases or decreases, the body maintains a nearly constant pH of 7.4. The body uses buffers to maintain this remarkable feat.

What is a buffer and how does it work?



A buffer consists of a weak acid and the salt of its conjugate base, or a weak base and the salt of its conjugate acid.

Examples:

HF + NaOH NaF + HOHw.a. c.b.



A buffer consists of a weak acid and the salt of its conjugate base, or a weak base and the salt of its conjugate acid.

Examples:

HF + NaOH NaF + HOHw.a. c.b.

NH3 + HCl NH4Cl w.b. c.a.

1.0 L

HF (g) NaF (s)

Buffer preparation: Add 0.10 mole HF (g) and NaF (s) to 1.0 L of water.

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1.0 L

HF (g) NaF (s)


HF (g) H+ + F-

NaF (s) Na+ + F-

H+ Na+ F-

HFlarge small

1.0 L


HF (g) H+ + F-

NaF (s) Na+ + F-

HFH+

Na+ F-

Now add the strong acid HCl

HCl

large small

1.0 L


HF (g) H+ + F-

NaF (s) Na+ + F-

HFH+

Na+ F-


HCl

HCl H+ + Cl-

H+ Cl-

What will the pH be if just water and no buffer?

Large small

1.0 L


HF (g) H+ + F-

NaF (s) Na+ + F-

H+

HF

Na+ F-


HCl

HCl H+ + Cl-

H+ Cl-

What will the pH be if just water and no buffer? pH = 1, dead if this is your blood.

Large small

1.0 L


HF (g) H+ + F-

NaF (s) Na+ + F-

H+

HF

Na+ F-


HCl

HCl H+ + Cl-

H+ Cl-


Large small

What removes the H+ to keep the pH near 7?

1.0 L


HF (g) H+ + F-

NaF (s) Na+ + F-

H+

HF

Na+ F-


HCl

HCl H+ + Cl-

H+ Cl-


Large small

What removes the H+ to keep the pH near 7? The conjugate base, F-

H+ + F- HF (a weak acid, low H+ )

1.0 L


HF (g) H+ + F-

NaF (s) Na+ + F-

H+

HF

Na+ F-

Now add the strong base NaOH

NaOH

Na+ OH-Large small

What will the pH be if just water and no buffer?

NaOH Na+ + OH-

1.0 L


HF (g) H+ + F-

NaF (s) Na+ + F-

H+

HF

Na+ F-

Now add the strong base NaOH

NaOH

Na+ OH-Large small

What will the pH be if just water and no buffer? PH = 13, dead again

NaOH Na+ + OH-

What removes the OH- to keep the pH near 7? The acid HF

HF + OH- F- + HOH

TitrationTitration is an experimental procedure to determine the concentration of an unknown acid or base.

The figure on the left shows the glassware for a titration experiment. A buret clamp holds the buret to a ring stand and below the buret is a flask containing the solution to be titrated, which includes an indicator. The purpose of the indicator is to indicate the point of neutralization by a color change.

The picture on the left shows the tip of a buret, with air bubble, which is not good, and also shows the stop-cock. Note the position of the stop-cock is in the “off” position. This picture shows the color of the phenolphthalein indicator at the end-point. In this experiment a 23.00 mL aliquot of 0.1000 M NaOH titrant is added to 5.00 mL of an unknown HCL solution. The acid solution in the beaker starts out clear and becomes pink when all of the HCL has been consumed.

NaOH + HCl NaCl + HOH

How can we calculate the concentration of acid in the beaker?

Titration


Normal procedure, yes, a conversion. Steps 1-4, again!

Titration



0.100 mole NaOHL NaOH solution




10-3 L solutionmL solution


Normal procedure, yes, a conversion. Steps 1-4,

again!



23.00 mL soln


Normal procedure, yes, a conversion. Steps 1-4,

again!



23.00 mL solnmole NaOHmole HCl




10-3 L solution

mL solution







10-3 L HCl soln.mL HCl soln.



0.100 mole NaOH

L NaOH solution

10-3 L solution

mL solution

23.00 mL soln

mole NaOH

mole HCl

10-3 L HCl soln.

mL HCl soln.

5.00 mL



0.100 mole NaOH

L NaOH solution

10-3 L solution

mL solution


10-3 L HCl soln.mL HCl soln.

5.00 mL=

0.460 M HCl

IndicatorsIndicators are weak organic (carbon containing) acids of

various colors depending on the formula of the acid.

Below is a generic acid.

HA H+ + A- colorless pink

1. Describe the color change when a strong acid is added?






Less pink






2. Describe the color change when a strong base is added?

Less pink


various colors depending on the formula of the acid. Below is

a generic acid.




Less pink

Darker pink







3. Describe the color change when the pH is lowered?

Less pink

Darker pink








Less pink

Darker pink

Less pink








4. Describe the color change when the pH is raised?

Less pink

Darker pink

Less pink








4. Describe the color change when the pH is raised?

Less pink

Darker pink

Less pink

Darker pink

Color versus pH of Many Different indicators

How can we make an indicator?

How can we make an indicator?

Step One

Red Cabbage

Step Two

Cook the Cabbage

Step Three

Filter the Juice

What color is the juice after filtering?

What color is the juice after filtering? The color of pH 6, 7, or

8

Colors of cabbage juice at various pH values

ACIDS BASES AND SALTS

The End Ch#14

chapter #14 acids, bases, and salts. acids, bases and salts topics the arrhenius theory the...

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