Acid-Base Chemistry





Acid Base Chemistry 3284
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Acids and bases have been known by their properties since the early days of experimental chemistry. The word "acid" comes from the Latin acidus , meaning "sour" or "tart," since water solutions of acids have a sour or tart taste. Lemons, grapefruit, and limes taste sour because they contain citric acid and ascorbic acid ( vitamin C). Another common acid is vinegar, which is the sour liquid produced when apple cider, grape juice, or other plant juices ferment beyond the formation of alcohol. Vinegar is a 5 percent water solution of acetic acid. Besides having a sour taste, acids react with active metals to give hydrogen, they change the colors of indicators (for example, litmus turns from blue to red), and they neutralize bases. Bases change the colors of indicators (litmus turns from red to blue) and they neutralize acids. Hence, bases are considered the chemical opposite of acids.

Most common acid-base reactions take place in water solutions (commonly referred to as aqueous solutions ). One of the earliest definitions of acids, advanced by the Swedish physicist and chemist Svante Arrhenius in 1887, stated that acid ionizes in aqueous solution to produce hydrogen ions (which are protons), H + , and anions ; and a base ionizes in aqueous solution to produce hydroxide ions (OH ) and cations. Later studies of aqueous solutions provided evidence of a small, positively charged hydrogen ion combining with a water molecule to form a hydrated proton, H + (H 2 O) or H 3 O + , which is called the hydronium ion. Often, the hydronium ion or hydrated proton is represented as H + ( aq ). Hydrogen chloride (HCl), a gas, is an acid because it dissolves in water to yield hydrogen ions and chloride ions. This water solution of HCl is referred to as hydrochloric acid.

         (1)

A typical base, according to the Arrhenius definition, is sodium hydroxide (NaOH). It dissolves in water to give sodium ions and hydroxide ions.

         (2)

Neutralization

In the reaction of an acid with a base in aqueous solution, the hydrogen ions of the acid react with the hydroxide ions of the base to give water. The second product is a salt, which is composed of the positive metal ion from the base and the negative ion from the acid. For example,

HCl ( aq ) + KOH ( aq ) → H 2 O ( l ) + KCl ( aq )          (3)

Since HCl ( aq ) and KOH ( aq ) are fully ionized in solution, the preceding equation can be written as

H + ( aq ) + Cl ( aq ) K + ( aq ) + OH ( aq ) → H 2 O ( l ) K + ( aq ) + Cl ( aq )          (4)

Ions common to both sides can be canceled to yield

H + ( aq ) + OH ( aq ) → H 2 O ( l )          (5)

This is referred to as the net ionic equation for the neutralization reaction. If H 3 O + is substituted for H + ( aq ), the neutralization equation becomes

H 3 O + ( aq ) + OH ( aq ) → 2 H 2 O ( l )          (6)

Strengths of Acids and Bases

The strength of an acid or base is determined by the extent of its ionization in aqueous solution. Strong acids, such as hydrochloric acid, are 100 percent ionized in aqueous solution, whereas weak acids, such as acetic acid, are less than 5 percent ionized. Experimentally, the extent of ionization is determined by measuring the electrical conductance of solutions. Strong acids and bases are strong electrolytes, and weak acids and bases are weak

Table 1. Common acids and bases
Table 1. Common acids and bases

COMMON ACIDS AND BASES
Strong Acids Strong Bases
HCl hydrochloric acid NaOH sodium hydroxide
HNO 3 nitric acid KOH potassium hydroxide
H 2 SO 4 sulfuric acid Ba(OH) 2 barium hydroxide
Weak Acids Weak Bases
CH 3 COOH acetic acid NH 3 ammonia
H 2 CO 3 carbonic acid CH 2 NH 2 methyl amine
H 3 PO 4 phosphoric acid

electrolytes. Table 1 lists some common acids and bases and indicates whether they are strong or weak.

For weak acids and bases, partial ionization is a dynamic equilibrium between unionized molecules and its ion, as indicated by the double arrow in equation (7). For example, acetic acid is only partially ionized in aqueous solution

CH 3 COOH ( aq ) ⇆ H + ( aq ) + CH 3 COO ( aq )          (7)

In acetic acid, hydrogen ions and acetate ions recombine to form acetic acid molecules. The double arrow signifies that at any given instant, less than 5 percent of acetic acid molecules dissociate into hydrogen ions and acetate ions, while the hydrogen ions and acetate ions recombine to form acetic acid molecules.

Ammonia (NH 3 ) is a weak base, and although it does not have OH ions in its formula, it produces the ion on reaction with water.

NH 3 ( aq ) + H 2 O ( l ) ⇆ NH 4 + ( aq ) + OH ( aq )          (8)

Brønsted-Lowry theory

A major problem with Arrhenius's acid-base theory is that some substances, like ammonia, produce basic solutions and react with acids, but do not contain hydroxide ions. In 1923 Johannes Brønsted, a Danish chemist, and Thomas Lowry, an English chemist, independently proposed a new way to define acids and bases. An acid donates hydrogen ions (also called a proton donor); a base accepts hydrogen ions (also called a proton acceptor). These definitions not only explain all the acids and bases covered by Arrhenius's theory, they also explain the basicity of ammonia and ions such as carbonate, CO 3 2− , and phosphate, PO 4 3 .

The Brønsted-Lowry theory includes water as a reactant and considers its acidity or basicity in the reaction. In the partial ionization of acetic acid, water is a base because it accepts the hydrogen ion to form hydronium ion.

         (9)

THOMAS M. LOWRY (1874–1936)

A meticulous experimenter, Thomas Lowry is best known for his conceptualization of acid–base chemistry. Studies of nitrogenous compounds led Lowry to question fundamental aspects of the role of hydrogen during acid–base reactions. Three months before Brønsted published his theory, Lowry released his own similar thoughts on proton acceptors and donors in print.

—Valerie Borek

In the reaction, a new acid and a new base are formed, which are called the conjugate acid and conjugate base, respectively. The hydronium ion, H 3 O + , is the conjugate acid of the base, H 2 O, and the acetate ion, CH 3 COO , is

Acid-Base Chemistry

CONJUGATE ACID-BASE PAIRS
Acid Conjugate Base
Strong acids H 2 SO 4 HSO 4 Weak bases
HCl Cl
H 3 O + H 2 O
HSO 4 SO 4
H 3 PO 4 H 2 PO 4
CH 3 COOH CH 3 COO
H 2 CO 3 HCO 3
H 2 PO 4 HPO 4 2−
NH 4 + NH 3
HCO 3 CO 3 2−
HPO 4 2− PO 4 3−
Weak acids H 2 O OH Strong bases

the conjugate base of acetic acid, CH 3 COOH. A pair of molecules or ions related to one another by the gain or loss of a single hydrogen ion is called a conjugate acid-base pair. In the reaction of ammonia, water is an acid because it donates a hydrogen ion to ammonia.

         (10)

This ability of water to donate or accept hydrogen ions, depending on whether it reacts with a base or an acid, is referred to as "amphiprotic." The conjugate acid-base pairs in this reaction are NH 3 /NH 4 + and H 2 O/OH .

The Brønsted-Lowry definitions also explain why carbonate salts such as sodium carbonate (washing soda) dissolve in water to give basic solutions. Carbonate ion removes a hydrogen ion from a water molecule, which leaves behind a hydroxide ion:

         (11)

In the preceding reaction, water and hydroxide ion are a conjugate acid-base pair, whereas carbonate ion and bicarbonate ion are a conjugate base-acid pair. Every Brønsted-Lowry acid has a conjugate base, and every Brønsted-Lowry base has a conjugate acid. Familiarity with conjugate acid-base pairs is important to understanding the relative strengths of acids and bases. Table 2 lists some conjugate acid-base pairs and their relative strengths. Strong acids have weak conjugate bases, and weak acids have strong conjugate bases.

Polyprotic Acids

Several common acids have more than one ionizable hydrogen ion (Table1). Each successive hydrogen ion in these polyprotic acids ionizes less readily. For example, sulfuric acid is a strong acid because of the complete ionization of the first hydrogen ion.

H 2 SO 4 ( aq ) + H 2 O ( l ) → H 3 O + ( aq ) + HSO 4 ( aq )          (12)

The HSO 4 also acts as an acid, but it is not 100 percent ionized, so HSO 4 is an acid of moderate strength. For example, sodium hydrogen sulfate is used to increase the acidity of swimming pools, whereas sodium carbonate is used to increase the basicity of swimming pools.

HSO 4 ( aq ) + H 2 O ( l ) ⇆ H 3 O + ( aq ) + SO 4 2− ( aq )          (13)

Phosphoric acid has three ionizable hydrogen ions. Each stepwise ionization of phosphoric acid occurs to a lesser extent than the one before it. Phosphoric acid is stronger than acetic acid because the first step ionizes to a greater extent than acetic acid.

H 3 PO 4 ( aq ) + H 2 O ( l ) ⇆ H 3 O + ( aq ) + H 2 PO 4 ( aq )          (14)

However, H 2 PO 4 is a weaker acid than acetic acid because the second ionization is much smaller (by a factor of 10 5 ) than the first step.

H 2 PO 4 ( aq ) + H 2 O ( l ) ⇆ H 3 O + ( aq ) + HPO 4 2− ( aq )          (15)

The third ionization is also much smaller than the second step (by a factor of 10 5 ).

HPO 4 2− ( aq ) + H 2 O ( l ) ⇆ H 3 O + ( aq ) + PO 4 3− ( aq )          (16)

The anions of phosphoric acid can also accept hydrogen ions and act as bases with a strong acid like hydrochloric acid. For example,

PO 4 3− ( aq ) + H 3 O + ( aq ) ⇆ HPO 4 2− ( aq ) + H 2 O ( l )          (17)

Ions such as H 2 PO 4 , HPO 4 2− , HCO 3 , and HSO 4 can act as an acid by donating a hydrogen ion or as a base by accepting a hydrogen ion. This ability to act as either an acid or a base is referred to as an amphoteric property.

The pH Scale

The Brønsted-Lowry acid-base definitions are based on the amphiprotic properties of water: Water is capable of acting as both a hydrogen ion donor and a hydrogen ion acceptor, depending on the acidic or basic properties of the dissolved substance (equations 9 and 10). Water can also act as a proton donor and proton acceptor towards itself. This is referred to as the autoionization of water.

H 2 O ( l ) + H 2 O ( l ) ⇆ H 3 O + ( aq ) + OH ( aq )          (18)

Pure water is neutral because it contains equal numbers of hydronium ions and hydroxide ions. However, pure water only slightly ionizes, about 1 in every 55,000,000 water molecules is ionized at any given time. The actual molar concentration of hydronium ions and hydroxide ions in pure water at 25°C is 1.0 × 10 −7 . The product of the molarity of the hydronium ions and hydroxide ions of pure water is (1.0 × 10 −7 ) (1.0 × 10 −7 ) = 1.0 × 10 −14 .

The value of 1.0 × 10 −14 is important to the study of aqueous solutions of acids and bases because it is a constant that is always the product of the molar concentration of H 3 O + and OH .

[H 3 O + ] [OH ] = 1.0 × 10 −14          (19)

If acid is added to pure water, the concentration of H 3 O + will be greater than 1.0 × 10 −7 , and then the concentration of OH will be less than 1.0 × 10 −7 . However, the product of the two must equal 1.0 × 10 −14 . This relationship is the basis for calculating the concentration of one of the two ions, hydronium or hydroxide, when the other one is known. For example, a 0.1 M solution of hydrochloric acid is 0.1 M in H 3 O + since hydrochloric

Table 3. pH of common substances.
Table 3. pH of common substances.

pH OF COMMON SUBSTANCES
pH [H 3 O + ], M Example
0 1.0 Battery acid, 1 M sulfuric acid
1 0.1 Stomach acid, 0.1 M hydrochloric acid
2 1 × 10 −2 Lemon juice
3 1 × 10 −3 Vinegar
4 1 × 10 −4 Soft drink
5 1 × 10 −5 Rain water
6 1 × 10 −6 Milk
7 1 × 10 −7 Pure water
8 1 × 10 −8 Baking soda, NaHCO 3
9 1 × 10 −9 Washing soda, Na 2 CO 3
10 1 × 10 −10 Milk of magnesia, Mg(OH) 2
11 1 × 10 −11 Aqueous household ammonia, NH 3
12 1 × 10 −12 Limewater, Ca(OH) 2
13 1 × 10 −13 Drano, 0.1 M NaOH
14 1 × 10 −14 Drano, 1.0 M NaOH

acid is fully ionized. From the equation, the molar concentration of OH is 10 −13 . For a 0.1 M solution of NaOH, the OH is 0.1 M , but the hydrogen ion concentration is 10 −13 . Hence, the value of the exponent for hydronium ion concentration goes from −1 in strong 0.1 M acid to −13 in strong 0.1 M base.

In 1909 the Danish biochemist S. P. L. Sørensen proposed that these exponents be used as a measure of acidity. He devised a scale that would be useful in testing the acidity of Danish beer. Sorensen's scale came to be known as the pH scale, from the French pouvoir hydrogene , which means hydrogen power. pH is defined as the negative logarithm (log) of the hydronium ion concentration.

pH = −log[H 3 O + ]          (20)

The brackets around hydronium ion mean moles per liter of hydronium ions.

The pH scale includes values between 0 and 14. The pH of pure water is 7 because [H 3 O + ] is 1.0 × 10 −7 . The pH of 0.1 M HCl is 1[−log 10 −1 = −(−1)]. The pH of 0.1 M NaOH is 13(−log 10 −13 ) = 13. The pH scale does not apply to concentrations greater than 1.0 M for a strong acid (pH = 0) or 1.0 M for a strong base (pH = 14).

For solutions in which [H 3 O + ] is not an exact power of 10 (0.1, 0.01,...), a calculator can be used to determine the logarithm. For example, if the [H 3 O + ] is 1.5 × 10 −3 M , the logarithm is −3 + log 1.5 = −3 + 0.18 = −2.82, and the pH is −(−2.82) or 2.82. Table 3 provides the pH values of some common solutions.

Acid-Base Indicators

Many natural substances are acid-base indicators. The most familiar one is litmus, an organic dye extracted from certain lichens. Litmus turns from blue to red in acidic solutions (< pH 7) and from red to blue in basic solutions (∼ pH 7). Some other natural indicators include red cabbage extract, blueberry juice, black tea, beet juice, rhubarb, and tomato leaves, and flowers such as the rose, daylily, blue iris, and purple dahlia. Red cabbage extract undergoes sharp changes of color at several pH values. The deep purple color of red cabbage leaves is caused by a mixture of water-soluble

Table 4. Common acid-base indicators.
Table 4. Common acid-base indicators.

COMMON INDICATORS
Indicator pH Range Color Change
Thymol blue 1.2 – 2.8 red → yellow
Methyl red 4.4 – 6.2 red → yellow
Litmus 5 – 8 red → blue
Bromothymol blue 6.2 – 7.6 yellow → blue
Phenolphthalein 8.0 – 10.0 colorless → pink

anthocyanins. Over the pH range of 2 to 12, these anthocyanins change from red (pH 2) to pink (pH 4) to purple (pH 6−7) to green (pH 10) to yellow (pH 12), which makes red cabbage extract a "universal indicator."

Acid-base indicators are weak acids and bases. A typical indicator will ionize in aqueous solution according to the equation

HIn ( aq ) + H 2 O ( l ) ⇆ H 3 O + ( aq ) + In ( aq )          (21)

The chemical species HIn and In are different colors. When the solution is acidic to the degree that the HIn species dominates, it will be the color of HIn. When the solution is more basic with In dominating, it will be the color of In . Some common indicators and the pH ranges for their color changes are listed in Table 4.

Buffer Solutions

Buffer solutions contain a base and an acid that can react with an added acid or base, respectively, and they maintain a pH very close to the original value. Buffers usually consist of approximately equal quantities of a weak acid and its conjugate base, or a weak base and its conjugate acid. For example, a buffer solution of acetic acid and its conjugate base, the acetate ion, can neutralize small amounts of a strong acid or strong base as follows:

CH 3 COOH ( aq ) + OH ( aq ) → CH 3 COO ( aq ) + H 2 O ( l )          (22)

CH 3 COO ( aq ) + H 3 O + ( aq ) → CH 3 COOH ( aq ) + H 2 O ( l )          (23)

As illustrated in equations (22) and (23), the addition of either a strong base or a strong acid produces one of the components of the buffer mixture and so the pH does not change. Buffers are limited in their buffer capacity, that is, the amount of a strong acid or strong base that can be added before the pH changes by 1 pH unit.

Buffers are very important to many industrial and natural processes. For example, controlling the pH of blood is essential to human health. The pH of blood is normally 7.40 ± 0.05, and good health depends on the ability of buffers to maintain the pH of blood within this narrow range. If the pH falls below 7.35, a condition known as acidosis occurs; increasing pH above 7.45 leads to alkalosis. Both these conditions can be life threatening. Two buffer systems, H 2 CO 3 /HCO 3 and H 2 PO 4 /HPO 4 2− , control the pH of the blood.

Lewis Acid-Base Theory

In the early 1930s Gilbert Lewis, an American chemist, proposed a more general acid-base theory that is based on sharing electron pairs rather than proton transfers. A Lewis acid is a substance that can accept a pair of electrons to form a new bond, and a Lewis base is a substance that can donate a pair of electrons to form a new bond. All Arrhenius and Brønsted-Lowry acids and bases are Lewis acids and bases. However, Lewis acid-base theory is more general because a Lewis base can donate an electron pair to something other than H + . For example, the gas phase reaction of NH 3 with BF 3 is a Lewis acid-base reaction.

         (24)

Solvent System Acid-Base Theory

Another acid-base theory that is useful for solvents other than water was postulated by American chemist Edward Franklin in 1905. It makes use of the autoionization of solvents, and defines an acid as a solute that produces the positively charged species of the solvent and a base as a solute that produces the negatively charged species of the solvent. In the case of the autoionization of water (equation 18) H 3 O + is the acid and OH is the base. For the nonaqueous solvent, liquid ammonia, the autoionization gives

NH 3 ( l ) + NH 3 ( l ) ⇆ NH 4 + + NH 2          (25)

so an acid in liquid ammonia is any solute that produces NH 4 + and a base in liquid ammonia is any solute that produces NH 2 . An example of an acid-base reaction in liquid ammonia is

         (26)

Note that liquid ammonia still falls within the Brønsted-Lowry definitions since NH 4 + is a proton donor and NH 2 is a proton acceptor.

Summary

The Brønsted-Lowry theory, which defines acids as proton donors and bases as proton acceptors, covers all acid-base reactions in aqueous solution. The strength of acids and bases is related to the percent of their ionization in water. Strong acids and bases are 100 percent ionized, whereas weak acids and bases are less than 5 percent ionized. There are a number of salts that have acidic or basic properties in solution. For example, baking soda, NaHCO 3 , can be used as an antacid because the bicarbonate ion, HCO 3 , is a strong enough conjugate base to combine with H 3 O + to give carbonic acid.

HCO 3 ( aq ) + H 3 O + ( aq ) → H 2 CO 3 ( aq ) + H 2 O ( l )          (27)

The pH scale is a convenient way to represent the acidity or basicity of dilute acid and base solutions. Pure water has a pH of 7; acidic solutions have pH values http:// 7 and basic solutions have pH values 7. Each change of one unit of pH is a tenfold change in acidity. Acid-base indicators, such as litmus and phenolphthalein, can be used to measure whether a solution is acidic or basic. A natural "universal indicator," red cabbage extract, can be used to determine the pH within 2 pH units. A buffer contains equal amounts of either a weak acid and its conjugate base or a weak base and its conjugate acid.

SEE ALSO Arrhenius, Svante ; Bases ; BrØnsted, Johannes Nicolaus ; Chemical Reactions ; Lewis, Gilbert N. ; Solution Chemistry .

Melvin D. Joesten

Bibliography

Atkins, Peter, and Jones, Loretta (1997). Chemistry: Molecules, Matter, Change , 3rd edition. New York: W. H. Freeman.

Joesten, Melvin D., and Wood, James L. (1996). The World of Chemistry , 2nd edition. Fort Worth, TX: Saunders College.

Moore, John W.; Stanitski, Conrad L.; Wood, James L.; Kotz, John C.; and Joesten, Melvin D. (1998). The Chemical World , 2nd edition. Philadelphia: Saunders.

Shakhashiri, Bassam Z. (1989). Chemical Demonstrations , Vol. 3. Madison: University of Wisconsin Press.

Internet Resources

More information available from http://www.visionlearning.com/library/index.htm

"CHEMystery: An Interactive Guide to Chemistry." Available from http://library.thinkquest.org/3659/acidbase/ .



Also read article about Acid-Base Chemistry from Wikipedia

User Contributions:

karlo
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Oct 11, 2007 @ 6:18 pm
where are the numbers 1 and 2? i can't see them men! please add more information about this topic! thanks! have a nice day!
Muwonge Henry
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Sep 5, 2008 @ 4:04 am
Good presentation and convincing. There are many things that I have learnt from the article and I wish to inform you that I will always have to access this article whenever I am in need of it. I will do this as I seek for missing information and add it if necessary.
Nokwanda
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Dec 3, 2009 @ 5:05 am
i just wanted to ask... why is it that a strong an a weak acid of the same concentration can produce the same amount of gas when reacted with a metal and also why does the weak acid take longer than he strong acid? thank you
Aleksandra Drizo
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Jul 28, 2010 @ 8:08 am
I use Ca rich material to purify wastewater, which generates effluents of highly elevated pH (~13).

I need to know what would be the best acid to add and what quantity to neutralize 1 m3 of my effluent?

Thank you very much,

Sincerely,

Ana
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Aug 23, 2010 @ 9:09 am
I think that this encyclopidia is best i get all the information that i want .so i think that everyone should try it.
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Sep 15, 2010 @ 8:08 am
Given 6 aqueous solutions of pH 2,1,9,5,6 and 11, How do we get two pairs of solutions that would react together to give a solution with a pH of 7?
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Oct 3, 2010 @ 4:04 am
Hello, can you please tell me what actually happens chemically when you have a Nitric acid and you add Sodium Bicarbonate or Calcium Carbonate to neutralize it.. Can you please explain to me what exactly is taking place in a chemical level?

I'm a Navy Explosive Odrnance Disposal tech, and I need to find out this answer...

Thanks,
EOD1 Hernandez
USN
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Nov 1, 2010 @ 2:14 pm
Hello,
can somebody help me :) i need to write the chemical equation showing how Milk of Magnesia neutralizes stomach acid???
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Nov 9, 2010 @ 7:19 pm
H3PO4 > < (equalibrium sign) H+ + H2PO4 -1

What will happen to the equilibrium if H+ is added?

What will happen to the equilibrium if OH is added?

What will happen to the equilibrium if H2O is added?
Corbin
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Nov 28, 2010 @ 5:17 pm
I don't understand the acids and bases thing. Is nitrogen an acid or a base?
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Nov 29, 2010 @ 5:05 am
please i need aresearch about the acidity of carboxylic acid and their adverties
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Apr 17, 2011 @ 7:19 pm
me and my friend are having a debate and was just wondering is there any way that you can change the concentration of acid without changing its chemical formula, haste of reply is paramount.
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Apr 22, 2011 @ 3:03 am
I want to know the chemical formula of a compound resulting from the development of phenolphthalen with ammonia solution ... Thank you
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May 3, 2011 @ 11:11 am
Pls. help! For each of the following salts, what are the formula of the base and acid from which the salt is formed, pls. classify whether the aqueous solution of the salt is acidic, basic or neutral. (assuming that the strenght of strong acids and bases are equal and the ionization constant of the weak acids and weak bases are equal.)
1. (NH4)2SO4
2. RbCN
3. MgSO4
4. KCL
5. KF
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May 4, 2011 @ 1:13 pm
chem acids and bases to assist you in the chemistry 110 class and possibly to help in test on tuesday may 10--hope this is found to be useful--take care and have a great day
giovanna
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May 16, 2011 @ 11:11 am
If they give me a compound, how do I know if it's an acid or a conjugate acid? And how can I know from what compounds the conjugate was created? PLEASE HELP.
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Jun 24, 2011 @ 1:01 am
I thanks you. I don't knew that who has provided this information but it was helpful to me. once again thanks a lot
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Aug 7, 2011 @ 3:03 am
why natural indicators such as colored flower extract changes it's color when mixed or combined with acid or base?
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Oct 10, 2011 @ 8:08 am
weak bases are used to neutralize the acidity of the stomach?


You have to give my question?
thanks =)
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Oct 19, 2011 @ 3:15 pm
IN WASTE WATER TREATMENT PLANT I WANT KNOW TO CALCULATE THE REQUIRED SOLUTION OF BASE SAY CAUSTIC TO NUETRALIZE THE EFFLUENT OF PH 6.2, TO GET AROUND 7.4.
Rahul
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Nov 14, 2011 @ 3:03 am
What is the bond formed in a salt? how it is made? does any bi products will be appeared while neutralization?
Varnika
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Dec 21, 2011 @ 8:08 am
How do we prepare a greeting card with BLACK GRAPE JUICE and BASE ?
Jule
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Jan 24, 2012 @ 11:11 am
In a recent lab:
part A. Measurement of pH of Some typical salt solutions
Record the pH of the solutions studied.
0.1 M NaCl= 7.03
0.1M NaC2H3O2= 8.73
0.1 M NH4NO3= 5.16
0.1 M Na2Co3= 11.27

Explain the observation of any pH in these solutions that is not within 1 pH unit of 7 by writing the net ionic equation for the reaction responsible for the pH change.
pete
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Feb 5, 2012 @ 1:13 pm
what causes the electrolyte in a battery to become fluffy white when I add sulfuric acid
pete
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Feb 6, 2012 @ 6:06 am
what causes the electrolyte in a battery to become fluffy white when I add sulfuric acid
SARPONG NELSON
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Feb 15, 2012 @ 2:02 am
How do you use the Arrhenius concept to explain the basicity of Na2C03
kamarudin
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Feb 23, 2012 @ 12:00 am
In presence of ascorbic acid nitrite is converted to nitrous oxide, NO. How to prevent the NO from being reversed back to nitrite
Lizzie Javapro
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Aug 27, 2013 @ 9:09 am
How do you use the Arrhenius concept to explain the presence of ascorbic acid nitrite in the basicity of Na2Co3

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