CAPITULO III DEFORMIDADES EN NAUPLIOS
3.2 Deformidades en antenas y el abdomen
When a diluting a solution, the number of moles of substance remains the same: only the volume of solvent changes. For that reason, the following formula can be used for diluting solutions:
V1⋅N1 = V2⋅N2
V = volumes of solutions (ml or L) N = normal concentrations (not molar!)
That formula is also used when titrating substances: at the endpoint of titration the number of equivalents of titrating solution is equal to that of titrated substance.
Example 1:
Prepare 500 ml of 0.01N solution, stating from 5N one.
500 ml⋅0.1N = V⋅5N V = 10 ml Take 10 ml of 5N solution and dilue with water to 500 ml.
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5.0 ml of AgNO3 0.1N have been used to titrate 1 ml of filtrate. What is the concentration of chloride ions?
5.0 ml⋅0.1N = 0.5 meq AgNO3 = meq Cl- (= mmol Cl-)
mgCl- in 1.0 ml of filtrate = mmol⋅MW = 0.5⋅35.45 = 17.72 mg mg/L Cl- = 17.72⋅1000 = 17720 mg/L
2.6 ACID AND BASES
In 1923, J. N. Bronsted in Denmark and T. M. Lowry in England independently, and almost simultaneously, proposed the modern "protonic" or Bronsted-Lowry theory of acid-base behavior. According to the Bronsted-Lowry concept, “an acid is any compound or ion which can give up a proton, while a base is any compound or ion which can accept a proton.”
A molecular species which can either accept or give up a proton is said to be amphiprotic. Thus the water molecule is amphiprotic, since it can give up a proton, H2O ⇒ H+ + OH-, to form the hydroxyl ion OH-. Alternatively, water can accept a proton to form the hydronium ion H3O+, according to the equation H+ + H2O ⇒ H3O
+
. The above two equations can be combined to give the dissociation equation for water: 2H2O ⇒ H3O+ + OH-.
The Bronsted-Lowry concept is an extension of the Arrhenius concept in that “bases, being sources of hydroxide, can accept protons; on the other hand acids, being sources of protons, can accept hydroxide.” Ammonia and amines will also accept protons to form the corresponding ammonium ions, so the existence of NH4OH is no longer necessary to explain the action of ammonia as a base. The Bronsted-Lowry concept also is useful in protonic solvents other than water, such as liquid ammonia or glacial acetic acid, where the Arrhenius concept is not useful. We will, however, generally confine our discussion to aqueous solutions because they are so much more important.
Lewis Acids and Bases
The basic principles of the Lewis theory of acid-base behavior were also set down in 1923, by the American physical chemist G. N. Lewis. The Lewis definitions of acids and bases are even more inclusive than the Bronsted definitions. The Lewis definitions are that “an acid is an electron-pair acceptor and a base is an electron-pair donor.”
H
N
H
H
Since a base like ammonia (above) has a lone pair of electrons, it can be considered to "donate" them to a proton in forming the conjugate acid NH4+. The Lewis definitions are used to explain the effect of compounds such as AlCl3, which acts as an acid in non-aqueous organic solvents, on organic reactions. In protonic solvents, however, they are far less useful than are the Bronsted definitions.
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2.6.1 pH
A Simple Definition
pH is a logarithmic measure of hydrogen ion concentration, originally defined by Danish
biochemist Søren Peter Lauritz Sørensen in 1909:
pH = - log[H+]
where log is a base-10 logarithm and [H+] is the concentration of hydrogen ions in moles per liter
of solution (M). According to the Compact Oxford English Dictionary, the "p" stands for the German word for "power", potenz, so pH is an abbreviation for "power of hydrogen".
The pH scale was defined because the enormous range of hydrogen ion concentrations found in aqueous solutions make using H+ molarity awkward. For example, in a typical acid-base titration, [H+] may vary from about 0.01M to 0.0000000000001M. It is easier to write "the pH varies from 2 to 13".
The hydrogen ion concentration in pure water around room temperature is about 1.0 × 10-7M. A pH of 7 is considered "neutral", because the concentration of hydrogen ions is exactly equal to the concentration of hydroxide (OH-) ions produced by dissociation of the water (Kw).
Kw = [H
+]⋅[OH-
] = 10-7M x 10-7M = 10-14 M2
Increasing the concentration of hydrogen ions above 1.0 × 10-7 M produces a solution with a pH of less than 7, and the solution is considered "acidic". Decreasing the concentration below 1.0 × 10-7 M produces a solution with a pH above 7, and the solution is considered "alkaline" or "basic". So:
• 0 < pH < 7: acidic solution
• pH = 7: neutral solution
• 7 < pH < 14: alkaline solution
pH is often used to compare solution acidities. For example, a solution of pH 1 is said to be 10 times more acidic than a solution of pH 2, because the hydrogen ion concentration at pH 1 is ten times the hydrogen ion concentration at pH 2. This is correct as long as the solutions being compared both use the same solvent. You can't use pH to compare the acidities in different solvents because the neutral pH is different for each solvent. For example, the concentration of hydrogen ions in pure ethanol is about 1.58 × 10-10 M, so ethanol is neutral at pH 9.8. A solution with a pH of 8 would be considered acidic in ethanol, but basic in water!
2.6.2 Ionization Constant
The inherent or intrinsic strength of an aqueous acid (or base) is its ability to remove a proton from (or donate a proton to) the solvent water or other ions and molecules in aqueous solutions. For quantitative comparisons between different aqueous acids or bases, this ability is compared with the ability of the solvent water itself. The reaction used is the reaction which corresponds to the ionization equilibrium whose equilibrium constant is called the ionization constant. In other words, “strengths of acids and bases are expressed quantitatively in terms of the values of their ionization constants.”
Aqueous ionization constants are quantitative measures of the tendency of the acid or base to either donate a proton, written as H3O+ or often simply H+, or accept a proton from water. The greater the value of the equilibrium constant, the greater the percentage of the acid or base that
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will be in ionized form. As a generalization, we can use the value of the ionization constant equal to 0.1 as the point of distinction between a strong acid and a weak acid. Thus a strong acid is one for which the value of the acid ionization constant Ka is large (>> 0.1) and a weak acid is one for which the value of the acid ionization constant Ka is small (<< 0.1). Likewise, a strong base is one for which the value of the base ionization constant Kb is large (>> 0.1) and a weak base is one for which the value of the base ionization constant Kb is small (<< 0.1).
HA D H+ + A- Ka =
[ ] [ ][ ]
HA A H+ ⋅ − BOH D OH- + B+ Kb =[ ] [ ][
]
BOH OH B+ ⋅ −AH = acid (i.e. HCl, H2SO4)
A- = conjugate base (i.e. Cl-, SO4-2) BOH = base (i.e. NaOH, NH4OH) B+ = conjugate acid (Na+, NH4+)
There are only a few common strong acids: HCl, HNO3, HClO4 and H2SO4. In the case of sulfuric acid, H2SO4, only the ionization of the proton from H2SO4 to give HSO4- is strong; the ionization of HSO4- to give SO42- is not strong, but weak. Common strong bases include NaOH and KOH. On reaction with water, CaO gives the strong base Ca(OH)2 and for that reason CaO is considered a strong base also, as are the oxides of sodium and potassium.
Using the values of the ionization constant as quantitative measures of acid strength is equivalent to the qualitative statement that a strong acid is an acid for which loss of the proton to water is essentially complete, while a weak acid is an acid for which loss of the proton is incomplete. Likewise, a strong base is a base for which acquisition of a proton from water is essentially complete while a weak base is a base for which acquisition of a proton is noticeably incomplete.
2.6.3 Acids and Bases
If we look at the two definitions (above) of an acid–base pair they are both telling us different things! The ionization constant tells us exactly how strong the acid is, the pH tells us the concentration of protons [H+] in solution! For Mud Engineers pH is the more important of the two, because by knowing the concentration of [H+] we can figure out the concentration of the alkaline ions in solution and base judgments of mud quality upon these results.
In general terms, acids in water solutions have the following properties. 1. Sour taste (not recommended).
2. The ability to make litmus dye turns red.
3. The ability to make other indicators change to characteristic colors. 4. The ability to react with and dissolve certain metals to form salts. 5. The ability to react with abase or alkaline to form salts
Acids are classified as strong or weak according to their ability to donate their proton or concentration of the hydrogen ions in solution. Sulfuric acid (H2SO4) is a strong acid (pH < 1), (higher concentration of protons), while carbonic acid (H2CO3) is a weak acid (pH = 3). Water is
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so weak that is classified as neutral (pH = 7). Strong acids are corrosive and dangerous to skin, eyes and mucous membranes.
Common organic acids have different nomenclature but all share a common acid functionality. • Carboxylic acids (has a COOH group)
• Fatty acids (has a COOH group) • Amino acids (has a COOH group)
• Dicarboxylic acids (contains 2 COOH groups) Some common organic acids
CH
3OH
O
H
OH
O
CO
2H
H
H
CO
2H
HO
H
H
CO
2H
Formic
Acetic
Citric
Inorganic acids or mineral acids, include: sulfuric (H2SO4), hydrochloric (HCl), hydrofluoric
(HF), nitric (HNO3) and phosphoric (H3PO4).
The disassociation reaction for organic and mineral acids is the same.
CH3 OH O Acetic Acid H+ + CH3 O- O Proton HCl H+ + Cl-
Hydrochloric acid Proton
In general terms, bases in water solutions have the following properties.
1. Bitter taste.
2. The ability to make litmus dye turns blue.
3. The ability to make other indicators turn characteristic colors. 4. The ability to react with acids to form salts.
Bases are classified as strong or weak according to their ability to accept a proton or decrease the concentration of the hydrogen ions in solution. Potassium hydroxide (KOH) is a strong base (pH = 14), while sodium bicarbonate (NaHCO3) is a weak base (pH=8.4). Basic solutions range in pH from 7.1 to 14. Like acids, strong bases are corrosive to skin, eyes and mucous membranes.
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