A.P. Chem Report on Acids and Bases

8/6/2019 A.P. Chem Report on Acids and Bases

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JOHN F. KENNEDY CATHOLIC HIGH SCHOOL

Report on Acids and Bases

Jonathan Mullen

6/2/2011


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Jonathan Mullen Report on Acids and Bases June 2, 2011

Acid and Base Theories

Although there are many theories regarding acids and bases, there are three theories

which are most commonly utilized. These are: The Arrhenius Theory, The Brnsted-Lowry

Theory, and the Lewis Theory.

The Arrhenius Theory of Acids and Bases was introduced in the late eighteenth century

A.D. by a Swedish scientist known as Svante Arrhenius. His theory basically stated that

substances which segregate in water (H2O) and create (yield) electrically charged molecules

or atoms (known as ions), (Hydrogen ion [H+] are considered acids). Arrhenius definition

of a base was fairly similar to that of an acid, except for the fact that bases ionize in water

(H2O) and yield hydroxide ions (OH-). It is important to note that a hydrogen ion can not

exist alone in an aqueous solution (water solution), instead it exists in a combined state

with water molecules, known as the hydronium ion (H3O+).

Figure 1): Water as both base and acid. One H2O acts as a base and gains H+ to become H3O+; the other H2O

acts as an acid and loses H+ to become OH-

The Brnsted-Lowry Theory of Acids and Bases was proposed by Johannes Nicolaus

Brnsted, along with Thomas Martin Lowry in the early twentieth century A.D. This theory

defines acids as a molecule or an ion which is able to lose, (donate) a hydrogen cation, which

is also known as a proton (H+). This theory also defines a base as a substance which is able to

accept (gain) a proton.


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Figure 2): Brnsted-Lowry Theory

The Lewis Theory of acids and bases defines an acid (typically referred to as a Lewis

acid when referenced to this theory) as any substance which accepts lone pair electrons. A

Lewis base is defined as any substance which donates (gives off) lone pair electrons. Lone

pair electrons are defined as a valence electron pair that is not considered to be bonding or

sharing with other atoms. They are found in the outer-most orbit (electron shell) of an atom.

Figure 3): Lewis diagram showing the formation of the ammonium ion.Unlike the Brnsted-Lowry Theory, the Arrhenius Theory is very specific. The Arrhenius

theory involves a limited number of acid-base reactions while The Brnsted-Lowry theory is

much more general. The Brnsted Lowry Theory involves acid-base reactions that would

not be considered as such by the Arrhenius theory. It is important to note that the Lewis

theory is extremely broad and a very large amount of reactions, including a large amount

of organic reactions, can be classified as a Lewis acid-base reaction.


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Comparison between the properties of acids and bases

Acids and bases both have different chemical and physical (structural) characteristics.

While acids taste sour, react with certain metals (such as Zn, Fe, ect) to produce hydrogen

gas. Acids react with limestone (CaCO3) to produce carbon dioxide (CO2) and also react

with bases to form salts and water, Bases taste bitter, react with oils and grease, and react

with acids to form salts and water.

Acids must have a proton that can be donated, Substances which do not contain

hydrogen can not be acids. Therefore substances such as N3O can not be or act as acids.

Bases must be able to accept a proton, therefore, all bases must have an unshared pair of

electrons (as discussed in the Brnsted-Lowry Theory above), and cant be positively

charged. It is important to note that the stronger the negative charge (of a base) the stronger

the base.

Figure 4): The Lewis structure of ammonia (NH3, a common base), is shown above; notice the unshared

electron pair on the nitrogen.

Applications ofAcids and Bases

Household use of acids and bases mainly revolve around cooking and cleaning. For

example, Boric acid is a weak acid used as antiseptic and cleanser. Sodium hypochlorite gets

our whites whiter and is commonly referred to as bleach. Acids as well as bases help make

soaps and unclog pipes. The most used home acids are acetic acid (CH3COOH) which is the

Unshared electron pair


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main component of vinegar and citric acid (C6H8O7) which is used as a preservative, and is

naturally found in nearly every juice drink.

Industrial use of acids and bases are involved in nearly every aspect of manufacturing,

and constructing. One of the most used acids around the world is sulfuric acid (H2SO4).

Besides its common presence in ore processing, fertilizer manufacturing, oil

refining, wastewater processing, and chemical synthesis, sulfuric acid is the main component

in lead-acid batteries. A commonly used and very abundant base would be limestone.

Limestone (lime) is found in cement, masonry products, and fertilizers. It can also be added

to chemicals to increase the pH of a given substance.

Figure 5): Battery Room for a UPS System. Because the Acid-Lead cavity is surrounded by H 2O, the hydrogen

from the water can escape so proper ventilation and H2 gas detectors are needed

The pH scale and methods ofdetermining the pH of a substance.

The pH scale measures how acidic or basic a substance is. The pH scale ranges from 0

(extremely acidic) to 14 (extremely basic). A pH of 7 is neutral. In other words, a pH reading

of less than 7 is acidic and a pH greater than 7 is basic.

Notice the hydrogendetection system


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Because the pH scale is logarithmic, each whole pH value below 7 is ten times more

acidic than the next higher value. For example, pH 3 is ten times more acidic than pH 4 and

100 times more acidic than pH 5. The same holds true for all pH values above 7 (neutral),

each of which is ten times more alkaline (basic) than the next lower whole number value. For

example, pH 10 is ten times more alkaline (basic) than pH 9 and 100 times more alkaline

than pH 8.

Figure 6): The pH Scale

pH can be solved mathematically. The formula used to solve for pH is: (In

moles/lieter). To better explain the process of determining the pH of a substance, please see

the below example:

At equilibrium, the concentration of H+

is 1.00 10-7

, so we can calculate the pH of

water at equilibrium as H = -log[H+]= -log[1.00 10

-7] = 7.00 (neutral pH).


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At times, pH can be solved indirectly by determining the pOH of a given substance. The

formula used for determining the pOH of a substance is OH = -log [OH-]. This can be better

explained through the following example:

What is the pOH of a solution that has a hydroxide ion concentration of 4.82 x 10-5

M?

pOH = - log [4.82 x 10-5

] = - ( - 4.32) = 4.32.

The equation that links pH with pOH is; pH + pOH = 14. So if we have a pOH of 4.32

(from the example above) we can determine that the pH is: 9.68.

When pH cant be determined mathematically, indicators can be used to determine the

pH of a given substance. A pH indicator is a chemical which is added to a solution to

determine its pH visually. In other words, a pH indicator is a chemical detector

for hydronium ions (H3O+) or hydrogen ions (H

+) which changes the color of a solution

dependent upon the pH of the given solution. See the below table for common indicators

used for determining the pH of a solution.


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Figure 7): pH determining indicators (Source: http://en.wikipedia.org/wiki/File:PH_indicators.jpg)

Importance of Ka

Ka is an abbreviated symbol which refers to an acid dissociation constant. An acid

dissociation constant is the measure of the strength of an acid in a solution. College Board

defines Ka as: the equilibrium constant for a chemical reaction known as dissociation in the

context of acid-base reactions.


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The equilibrium mentioned above can be written symbolically as HA A

+ H+

(HA is

a generic acid that dissociates by splitting into A). A

-is known as the conjugate base of an

acid.(Note: H+

is the proton). A, A

and H+

are considered to be in a state of equilibrium

when their concentrations do not change over time.

More commonly used is the logarithmic function of Ka, which is demonstrated by the

below example.

The pKa is calculated using the expression:

pKa = - log (Ka)

where "Ka" is the equilibrium constant for the ionization of the acid.

Example: What is the pKa of acetic acid, if Ka for acetic acid is 1.78 x 10-5

?

pKa = - log (1.78 x 10-5

) = - ( - 4.75) = 4.75

It is important to know that pKa and Ka are not the same, however Ka can be determined

if there is a value for pKa.The Ka for an acid is calculated from the pKa by performing the

reverse of the mathematical operation used to find pKa.

Ka = 10-pKa

or Ka = antilog ( - pKa)

Example: Calculate the value of the ionization constant for the ammonium ion, Ka, if the pKa is

9.74.

9.74 = - log (Ka)-9.74 = log (Ka)

Ka = 10-9.74

= 1.82 x 10-1


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Strong Acids and WeakAcidsIt is easier to understand the concept of strong and weak acids when used in reference to

the Brnsted-Lowry Theory. As explained earlier, when an acid is dissolved in water (H2O),

a proton (hydrogen ion) is transferred to a water molecule to produce a hydroxonium ion and

a negative ion depending on what acid you are starting from. This can be written out as:

The reactions are reversible, but in a few cases, the acid is good at giving away hydrogen

ions that we can think of the reaction as being one-way.

Figure 8): At any one time, "virtually" 100% of the hydrogen chloride will have reacted to produce

hydroxonium ions and chloride ions. Hydrogen chloride is described as a strong acid.

In other words, A strong acid is one which is virtually 100% ionized in solution. (Note

that pH can aid in determining whether an acid is considered strong or weak. The lower the

pH, the higher the concentration of hydrogen ions in the solution.)

Unlike strong acids, weak acids are ones which do not ionize completely when dissolved

in water. This is best explained through the below example:

Ethanoic acid is a weak acid. It reacts with water and produces hydroxonium ions and

ethanoate ions, but the back reaction is more successful than the forward one. The ions react

easily to reform the acid as well as the water.


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Figure 9): Weak Acid Example.

There are several factors that affect acid strength, and can assist in determining an acids

strength.

1) Acid strength determined by H-A bond: As the strength of bond decreases, acid strength

increases.

2) As we go across periodic table from left to right, electro-negativity increases along with acid

strength.

3) For HOZ acids, high electro-negativity for Z leads to a stronger acid; Z withdraws

(pulls) electrons from O and this makes it easier to transfer protons.

4) The stronger acid typically has the greater number of oxygen atoms

Importance of Kb

Kb is the Base Dissociation Constant. weak bases (B), when placed into water, also

establish an equilibrium system much like weak acids (as described above):


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Similar to that of Ka, The equilibrium constant expression is called the weakbase dissociation

constant,Kb, and has the form:

The same description that was used regarding weak acids (above) is also true here: [HB+]

= [OH

]; [HB] Minitially; the numerator can be represented as [OH

]2; and knowing the initial

molarity andKb of the weak base, the [OH

] can easily be calculated. And if the

initial molarity and [OH

] are known,Kb can be calculated.

To better understand the base dissociation constant, lets propose that a 0.500 M solution

of ammonia has a pH of 11.48. How would we solve for Kb?


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The relationship between Ka and Kb

There is a simple relationship between Ka and Kb, the equilibrium constants for a weak

acid and a weak base. The rule for combining equations and equilibrium constants states that

when you add two chemical equations the resulting equilibrium constant K for the reaction is

the product of the first two. (Parts of this paragraph are from

http://www.cartage.org.lb/en/themes/sciences/Chemistry/Inorganicchemistry/AcidsBases/Ac

idsbasesindex/Relationship.htm)

Conjugate acids andconjugate bases

According to the Brnsted-Lowry theory of acids and bases, an acid is a proton donor and

a base is a proton acceptor. Therefore, once, an acid has given off a proton, the remaining

part can act as a proton acceptor, and therefore a base. Therefore, an acid and a base are

closely related to one another. This can be understood best by the following example:

For example:

NH3 + H2O =NH4+ + OH-

HAc = H+

+ Ac-

This shows thatNH4+

andNH3 are a pair of conjugate acids and bases, the same holds

true forHAc and Ac-. (Note: Colors used to distinguish acids between bases and are depicted on

the chart below)


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Figure 10): Various Conjugate acid/base pairs

Sources:

1) " B r n s t e d L o w r y a c i d b a s e t h e o r y . " W i k i p e d i a . W i k i m e d i a , 1 8 M A Y 2 0 1 1 . W e b .M A Y 2 0 1 1 .

h t t p : / / e n . w i k i p ed i a . o r g / w i k i / Br % C 3 % B 8 n s t e d % E 2 % 8 0% 9 3 L o w r y _a c i d % E 2 % 8 0 % 9

3 b a s e _ t h e o r y

2 ) "Arrhenius theory."Encyclopdia Britannica. Encyclopdia Britannica Online. Encyclopdia Britannica,

2011. Web. 01 Jun. 2011.

3 ) "L ewi s Ac ids and Bases . " W i k i p e d i a . W i k i m e d i a , 2 0 M a y 2 0 1 1 . W e b . 1 J u n 2 0 1 1 .

< h t t p : / / e n . w i k i p ed i a . o r g / w i k i / L e wi s _ a c i d s _ a n d _ b a s e s> .

4 ) " R e l a t i o n s h i p b e t w e e n K a a n d K b a n d K c . " C a r t a g e . G a r t a g e G r o u p L t d . , n . d .

W e b . 3 1 M A Y 2 0 1 1 .

h t t p : / / w w w. c a r t a g e . o r g . l b / e n / t h e m es / s c i e n c e s / C h e m i s t r y/ I n o r g a n i c c h e m i s t r y/ A c i ds B a s e s / A c i d s b a s e s i n d e x / R e l a t i o n s h i p . h t m

5 ) C l a r k , J i m . " s t r o n g a n d w e a k a c i d s . " C h e m g u i d e . G a r t a g e G r o u p L t d . , 2 0 0 2 . W e b .

1 J u n 2 0 1 1 . < h t t p : / / w w w . c h e m g u i d e. c o . u k / p h y s i c a l / a c i d b a s e e q i a / a ci d s . h t m l> .

6 ) " C a l c u l a t i n g _ p H a n d p O H. " P u r d u e C o l l e g e o f S c i e n c e s . P u r d u e , n . d . W e b . 1 J u n

2 0 1 1 .


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h t t p : / / w w w. c h e m . p u r d u e . e d u / g c h e l p/ h o w t o s o l v e i t / E q u i l i br i u m / C a l c u l a t i n g _ p Ha n d

pOH. h tm

7 ) " v C h e m b o o k . " pH Sca l e . W e b . 1 J u n 2 0 1 1 .

< h t t p : / / w w w. e l m h u r s t . e d u / ~ c h m / v ch e m b o o k / 1 8 4 p h . h t m l> .

8 ) " M o l e c u l a r S t r u c t u r e o f B a s e s . " C h e m P a g e s N e t o r i a l s . Un i v e r s i t y o f W i s c o n s i n ,n . d . W e b . 1 J u n 2 0 1 1 .h t t p s : / / w w w. c h e m . w i s c . e d u / d e p t fi l e s / g e n c h e m / n e t o r i a l / R Ot t o s e n / t u t o r i a l / m o d ul e s

/ a c i d _ b a s e / 0 2 m o l e c u l a r / m o l e c u l a r 3 . h t m

9 ) H a r d i n g e r , S t e v e . " A c i d s a n d B a s e s : M o l e c u l a r S t r u c t u r e a n d A c i d i t y. " U C L A

C h e m i s t r y a n d B i o c h e m i s t r y . Un i v e r s i t y o f C a l i f o r n i a L o s A n g l e s , 0 6 M A Y 2 0 0 9 .

W e b . 2 9 M A Y 2 0 1 1 .

< h t t p : / / w w w. c h e m . u c l a . e d u / h a r d i n g / t u t or i a l s / a c i d s _ a n d _ ba s e s / m o l _ s t r . p d f > .

1 0 ) " C h a p t e r 1 2 - A c i d s a n d B a s e s . " M o d e s t o J u n i o r C o l l e g e . M o d e s t o J u n i o r C o l l e g ,

n . d . W e b . 1 J u n 2 0 1 1 .

h t t p : / / v i r t u a l . y os e m i t e . c c . c a . u s/ l m a k i / C h e m 1 4 2 / c h a p _ o ut l i n e s / c h a p t e r 1 2 . h t m

1 1 ) " L o n e P a i r . " W i k i p e d i a . W i k i m e d i a , 2 A p r i l 2 0 1 1 . W e b .

< h t t p : / / e n . w i k i p ed i a . o r g / w i k i / L on e _ p a i r _ e l e c t r o n s> .

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A.P. Chem Report on Acids and Bases