Sunday, January 2, 2011

ACIDS AND BASES

HISTORY
For thousands of years people have known that vinegar, lemon juice and many other foods taste sour. However, it was not until a few hundred years ago that it was discovered why these things taste sour - because they are all acids. The term acid, in fact, comes from the Latin term acere, which means "sour". While there are many slightly different definitions of acids and bases, in this lesson we will introduce the fundamentals of acid/base chemistry.

In the seventeenth century, the Irish writer and amateur chemist Robert Boyle first labeled substances as either acids or bases (he called bases alkalies) according to the following characteristics:

Acids taste sour, are corrosive to metals, change litmus (a dye extracted from lichens) red, and become lessacidic when mixed with bases.

Bases feel slippery, change litmus blue, and become less basic when mixed with acids.
While Boyle and others tried to explain why acids and bases behave the way they do, the first reasonable definition of acids and bases would not be proposed until 200 years later.

In the late 1800s, the Swedish scientist Svante Arrhenius proposed that water can dissolve many compounds by separating them into their individual ions. Arrhenius suggested that acids are compounds that contain hydrogen and can dissolve in water to release hydrogen ions into solution.



Arrhenius defined bases as substances that dissolve in water to release hydroxide ions (OH-) into solutions.





The Arrhenius definition of acids and bases explains a number of things. Arrhenius's theory explains why all acids have similar properties to each other (and, conversely, why all bases are similar): because all acids release H+ into solution (and all bases release OH-). The Arrhenius definition also explains Boyle's observation that acids and bases counteract each other. This idea, that a base can make an acid weaker, and vice versa, is called neutralization.

Neutralization
As you can see from the equations, acids release H+ into solution and bases release OH-. If we were to mix an acid and base together, the H+ ion would combine with the OH- ion to make the molecule H2O, or plain water:
H+(aq)
+  
OH-(aq)
 =
H2O
The neutralization reaction of an acid with a base will always produce water and a salt, as shown below:
Acid
Base
Water
Salt
HCl
 + 
NaOH
 =
H2O
 + 
NaCl
HBr
 + 
KOH
 =
H2O
 + 
KBr
Though Arrhenius helped explain the fundamentals of acid/base chemistry, unfortunately his theories have limits. For example, the Arrhenius definition does not explain why some substances, such as common baking soda (NaHCO3), can act like a base even though they do not contain hydroxide ions.
In 1923, the Danish scientist Johannes Brønsted and the Englishman Thomas Lowry published independent yet similar papers that refined Arrhenius' theory .  In Brønsted's words, "... acids and bases are substances thatare capable of splitting off or taking up hydrogen ions, respectively."  The Brønsted-Lowry definition broadened the Arrhenius concept of acids and bases.  

The Brønsted-Lowry definition of acid is very similar to the Arrhenius definition, any substance that can donate a hydrogen ion is an acid (under the Brønsted definition, acids are often referred to as proton donors because an H+ ion, hydrogen minus its electron, is simply a proton). 

The Brønsted definition of bases is, however, quite different from the Arrhenius definition.  The Brønsted bases is defined as any substance that can accept a hydrogen ion.  In essence, a base is the opposite of an acid.  NaOH and KOH, as we saw above, would still be considered bases because they can accept an H+ from an acid to form water.  However, the Brønsted-Lowry definition also explains why substances that do not contain OH- can act like bases.  Baking soda (NaHCO3), for example, acts like a base by accepting a hydrogen ion from an acid as illustrated below:
Acid
Base


Salt
HCl
 + 
NaHCO3
 =
H2CO3
 + 
NaCl
In this example, the carbonic acid formed (H2CO3) undergoes rapid decomposition to water and gaseous carbon dioxide, and so the solution bubbles as CO2 gas is released.

pH
Under the Brønsted-Lowry definition, both acids and bases are related to the concentration of hydrogen ions present.  Acids increase the concentration of hydrogen ions, while bases decrease the concentration of hydrogen ions (by accepting them).  The acidity or basicity of something, therefore, can be measured by its hydrogen ion concentration.

In 1909, the Danish biochemist Sören Sörensen invented the pH scale for measuring acidity.  The pH scale is described by the formula:

pH = -log [H+]
Note: concentration is commonly abbreviated by using square brackets, thus [H+] = hydrogen ion concentration.  When measuring pH, [H+] is in units of moles of H+ per liter of solution.

For example, a solution with [H+] = 1 x 10-7 moles/liter has a pH equal to 7 (a simpler way to think about pH is that it equals the exponent on the H+ concentration, ignoring the minus sign). The pH scale ranges from 0 to 14. Substances with a pH between 0 and less than 7 are acids (pH and [H+] are inversely related - lower p H means higher [H+]). Substances with a pH greater than 7 and up to 14 are bases (higher pH means lower [H+]). Right in the middle, at pH = 7, are neutral substances, for example, pure water. The relationship between [H+] and pH is shown in the table below alongside some common examples of acids and bases in everyday life.
[H+]
pH
Example
Acids
1 X 100
0
HCl
1 x 10-1
1
Stomach acid
1 x 10-2
2
Lemon juice
1 x 10-3
3
Vinegar
1 x 10-4
4
Soda
1 x 10-5
5
Rainwater
1 x 10-6
6
Milk
Neutral
1 x 10-7
7
Pure water
Bases
1 x 10-8
8
Egg whites
1 x 10-9
9
Baking soda
1 x 10-10
10
Tums® antacid
1 x 10-11
11
Ammonia
1 x 10-12
12
Mineral lime - Ca(OH)2
1 x 10-13
13
Drano®
1 x 10-14
14
NaOH


Acids:
acid, is a substance which reacts with a base. Commonly, acids can be identified as tasting sour, reacting with metals such as calcium, and bases like sodium carbonate. Aqueous acids have pHs of less than 7, and turn blue litmus paper red. Chemicals or substances having the property of an acid are said to be acidic.

There are three common definitions for acids: the Arrhenius definition, the Brønsted-Lowry definition, and the Lewis definition. The Arrhenius definition states that acids are substances which increase the concentration of hydronium ions (H3O+) in solution. The Brønsted-Lowry definition is an expansion: an acid is a substance which can act as a proton donor. Most acids encountered in everyday life are aqueous solutions, or can be dissolved in water, and these two definitions are most relevant.


PROPERTIES OF ACIDS
  • Corrosive ('burns' your skin)
  • Sour taste (e.g. lemons, vinegar)
  • Contains hydrogen ions (H+) when dissolved in water
  • Has a pH less than 7
  • Turns blue litmus paper to a red color
  • Reacts with bases to form salt and water
  • Reacts with metals to form hydrogen gas
  • Reacts with carbonates to form carbon dioxide, water and a salt
EXAMPLES OF ACIDS
  • Hydrochloric acid (HCl) in gastric juice
  • Sulphuric acid (H2SO4)
  • Nitric acid (HNO3)
  • Carbonic acid in softdrink (H2CO3)
  • Uric acid in urine
  • Ascorbic acid (Vitamin C) in fruit
  • Citric acid in oranges and lemons
  • Acetic acid in vinegar
  • Tannic acid (in tea and wine)
  • Tartaric acid (in grapes)


    Base:
    A base in chemistry is a substance that can accept hydrogen ions or more generally, donate electron pairs. The Brønsted-Lowry theory defines bases as proton (hydrogen ion) acceptors, while the more general Lewis theory defines bases as electron pair donors, allowing other Lewis acids than protons to be included

    Bases can be thought of as the chemical opposite of acids. A reaction between an acid and base is called neutralization. Bases and acids are seen as opposites because the effect of an acid is to increase the hydronium ion (H3O+) concentration in water, whereas bases reduce this concentration. Bases and acids are typically found in aqueous solution forms. Aqueous solutions of bases react with aqueous solutions of acids to produce water and salts in aqueous solutions in which the salts separate into their component ions. If the aqueous solution is a saturated solution with respect to a given salt solute any additional such salt present in the solution will result in formation of a precipitate of the salt.


    PROPERTIES OF BASES AND ALKALIS
    • Corrosive ('burns' your skin)
    • Soapy feel
    • Has a pH more than 7
    • Turns red litmus paper to a blue colour
    • Many alkalis (soluble bases) contain hydroxyl ions (OH-)
    • Reacts with acids to form salt and water

    EXAMPLES OF BASES AND ALKALIS
    • Sodium hydroxide (NaOH) or caustic soda
    • Calcium hydroxide ( Ca(OH)2 ) or limewater
    • Ammonium hydroxide (NH4OH) or ammonia water
    • Magnesium hydroxide ( Mg(OH)2 ) or milk of magnesia
    • Many bleaches, soaps, toothpastes and cleaning agents




    Common Acids
    pH
    Common Bases
    pH
    Hydrochloric acid
    Sulphuric acid
    Stomach juice
    Lemons
    Vinegar
    Apples
    Oranges
    Grapes
    Sour milk
    White bread
    Fresh milk    		 0.1
    0.3
    1-3
    2.3
    2.9
    3.1
    3.5
    4
    4.4
    5.5
    6.5    		 Human saliva
    Distilled water
    Blood plasma
    Eggs
    Seawater
    Borax
    Milk of magnesia
    Ammonia water
    Limewater
    Caustic soda    		 6-8
    7
    7.4
    7.8
    7.9
    9.2
    10.5
    11.6
    12.4
    14

    INDICATORS

    An indicator, when added to an acid, a neutral substance or a base, will change different colours.

    INDICATOR COLOUR IN ACID COLOUR IN NEUTRAL SOLUTION COLOUR IN BASE
    Litmus red purple blue
    Bromothymol Blue yellow blue blue
    Phenolphthalein clear clear purple
    Universal Indicator red Yellow-green purple


    Sources:
    http://www.chem4kids.com
    http://www.visionlearning.com
    http://www.qldscienceteachers.tripod.com


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    Kyra Yu
    Lindsey Zuleta
    Erika Villa ignacio
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