Volume reading – Average volume reading
--------------------------------------------------- X 1000 ≤ 3
Average volume reading
Data / Results:
Titration of NaOH with HCl
Titration of NaOH with H2SO4
Analysis & Calculations:
- Titration of NaOH with HCL
(a) Mean Titre
Average = X1 + X2 + X3
---------------
3
= 11.80 + 12.00 + 11.80
---------------------------
3
= 11.87cm³
(b) Determination of accuracy
X1 = 11.80 – 11.87
----------------- X 1000 ≤ 3
11.87
= | -5.89 |
= 5.89cm³
X2 = 12.00 – 11.87
----------------- X 1000 ≤ 3
11.87
= 2.53cm³
X3 = 11.80 – 11.87
----------------- X 1000 ≤ 3
11.87
= | -5.89|
= 5.89cm³
(c) Concentration of base (diluted solution)
M1V1 = M2V2
M1(11.87) = (0.01)(25)
M1 = 0.25
-------
11.87
= 0.0211M
(d) Concentration of base (original solution)
M1V1 = M2V2
M1(50) = (0.0211)(250)
M1 = (0.0211)(250)
-----------------
50
= 0.1055M
Discussions:
Acid
An acid (often represented by the generic formula AH) is typically a water-soluble, sour-tasting chemical compound. In common usage an acid is any substance that, when dissolved in water, gives a solution with a pH of less than 7. In general scientific usage an acid is a molecule or ion that is able to give up a proton (H+ ion) to a base, or accept an unshared pair of electrons from a base. An acid reacts with a base in a neutralization reaction to form a salt.
Chemical characteristics
In water the following reaction occurs between an acid (AH) and water, which acts as a base:
The acidity constant (or acid dissociation constant) is the equilibrium constant for the reaction of AH with water:
Strong acids have large Ka values (i.e. the reaction equilibrium lies far to the right, lots of H3O+ present; the acid is almost completely dissociated). For example, the Ka value for hydrochloric acid (HCl) is 107.
Weak acids have small Ka values (i.e. at equilibrium significant amounts of AH and A- exist together in solution; modest levels of H3O+ are present; the acid is only partially dissociated). For example, the Ka value for acetic acid is 1.8 x 10-5.
Strong acids include the hydrohalic acids - HCl, HBr, and HI. (However, hydrofluoric acid, HF, is relatively weak.) Oxyacids, which tend to contain central atoms in high oxidation states surrounded by oxygen, are also quite strong and include HNO3, H2SO4, HClO4. Most organic acids are weak acids.
A few clarifications:
The terms "hydrogen ion" and "proton" are used interchangebly; both refer to H+.
In chemical equations H+ is often written, although in water it will actually be H3O+.
The strength of an acid is measured by its Ka value. pH measures how many hydrogen ions are present, which depends on both the type of acid (or base) and how much is there.
Acid strength is also defined by pKa= - log(Ka).
Number of acid dissociations
Some acid molecules are able give up more than one H+ ion (proton). Those acids which can give up only one H+ ion per molecule are called monoprotic acids, those acid molecules that can give up two H+ ions are diprotic acids, those that can give up three are triprotic acids, etc. A monoprotic acid can undergo one dissociation (sometimes called ionization) as follows and simply has one acid dissociation constant as shown above:
AH + H2O → A- + H3O+ Ka
A diprotic acid (here symbolized by AH2) can undergo one or two dissociations depending on the conditions (namely pH). Each dissociation has its own dissociation constant, Ka1 and Ka2.
AH2 + H2O → AH- + H3O+ Ka1
AH- + H2O → A-2 + H3O+ Ka2
The first dissociation constant is typically greater than the second; i. e. Ka1 > Ka2 . For example, sulfuric acid (H2SO4) can give up one H+ to form the singly charged bisulfate anion (HSO4-), for which Ka1 is very large; then it can give up a second H+ to form the doubly charged sulfate anion (SO4-2) where the Ka2 is intermediate strength. The large Ka1 for the first dissociation makes sulfuric a strong acid. Similarly, the weak unstable carbonic acid (H2CO3) can lose one H+ to form a singly charged bicarbonate anion (HCO3-) and lose a second to form a doubly charged carbonate anion (CO3-2). Both Ka values are small, but Ka1 > Ka2 .
Analogously, a triprotic acid (AH3) can undergo one, two, or three dissociations and has three dissociation constants, where Ka1 > Ka2 > Ka3 .
AH3 + H2O → AH2- + H3O+ Ka1
AH2- + H2O → AH-2 + H3O+ Ka2
AH-2 + H2O → A-3 + H3O+ Ka3
An inorganic example of a triprotic acid is orthophosphoric acid (H3PO4), usually just called phosphoric acid. All three of its H atoms can be successively lost as H+ (or H3O+ in water) to yield H2PO4-, then HPO4-2, and finally PO4-3 , the triply charged orthophosphate ion, usually just called phosphate. An organic example of a triprotic acid is citric acid, which can successively lose three H+ ions to finally form the triply charged citrate ion. Even though the positions of the H atoms on the original molecule may be equivalent, the successive Ka values will differ since it is energetically less favorable to lose a positive H+ if the ion is more negatively charged.
Acids are very dangerous. They are extremely reactive with metals and strong acids can give very serious burns just by touching them for an instant. If contact with an acid should occur seek medical attention immediately.
Different definitions of acid/base
The word acid comes from the Latin acidus meaning sour. In chemistry the term acid has a more specific meaning.
The Swedish chemist Svante Arrhenius defined an acid to be a substance that gives up hydrogen ions (H+) when dissolved in water (the product of the solution, H2O + H+, is called a hydronium ion, H3O+), while bases are substances that give up hydroxide ions (OH-). This definition limits acids and bases to substances that can dissolve in water. Later on, Brønsted and Lowry defined an acid to be a proton donor and a base to be a proton acceptor. In this definition, even substances that are insoluble in water can be acids and bases. The most general definition of acids and bases is the Lewis definition, given by the American chemist Gilbert N. Lewis. Lewis theory defines a "Lewis acid" as an electron-pair acceptor and a "Lewis base" as an electron-pair donor. It can include acids that do not contain any hydrogen atoms, such as iron(III) chloride. Acid/base systems are different from redox reactions in that there is no change in oxidation state. The Lewis definition can also be explained with molecular orbital theory. In general an acid can receive an electron pair in its lowest unoccupied orbital (LUMO) from the highest occupied orbital (HOMO) of a base. That is, the HOMO from the base and the LUMO from the acid combine to a bonding molecular orbital.
The Brønsted-Lowry definition, where an acid is treated as a proton donor, is sufficient for many situations. In this case, the proton (H+) is the actual acid and the acidity of the proton-donating-compound, such as an organic acid, is determined by its stability when it donates protons to the solution it is embedded in. So if the organic acid likes letting protons go, it has high acidity because it donates protons with empty molecular orbitals to the solution. This is how organic acids such as carboxylic acids work, here the Brønsted definition is nice for calculations while the Lewis definition is good for understanding.
Acid number
This is used to quantify the amount of acid present, for example in a sample of biodiesel. It is the quantity of base, expressed in milligrams of potassium hydroxide, that is required to neutralize the acidic constituents in 1 g of sample.
AN = (Veq-beq)×N×56.1/Woil.
Veq is the amount of titrant (ml) consumed by the crude oil sample and 1ml spiking solution at the equivalent point, and beqbeq is the amount of titrant (ml) consumed by 1ml spiking solution at the equivalent point.
The molarity concentration of titrant (N) is calculated as such: N = 1000×WKHP/(204.23×Veq).
In which, WKHP is the amount (g) of KHP in 50ml of KHP standard solution, and Veq is the amount of titrant (ml) consumed by 50ml KHP standard solution at the equivalent point.
Acid number (mgKOH/g oil) for biodiesel is preferred to be lower than 3.
Neutralization
Neutralization is a type of reaction between an acid and a base. The products include a salt and water. So, it is also called a water forming reaction
Example:
This type of reaction forms the basis of titration methods for analysing acids, where a pH indicator shows the point of neutralization.
Naming acids
Acids are named according to the ending of their anion. That ionic ending is dropped and replaced with a new suffix according to the table below. For example, HCl has chloride as its anion, so the -ide suffix makes it takes the form hydrochloric acid.
Base
The common (Arrhenius) definition of a base is a chemical compound that either donates hydroxide ions or absorbs hydrogen ions when dissolved in water. Bases and acids are referred to as opposites because the effect of an acid is to increase the hydronium ion concentration in water, whereas bases reduce this concentration. Arrhenius bases are water-soluble and always have a pH greater than 7 in solution.
Common bases
- Baking soda (sodium hydrogen carbonate)
- Sodium carbonate
- Ammonia and amines
- Pyridine and other basic aromatic rings
- Metal hydroxides like sodium hydroxide or potassium hydroxide
- Many metal oxides form basic hydroxides with water (anhydrides)
Bases and pH
The pH of (impure) water is a measure of its acidity. In pure water, about one in ten million molecules dissociate into hydrogen ions (H+) and hydroxide ions (OH−), according to the equation
The concentration (in mole/liter) of the ions is indicated as [H+] and [OH−]; their product is the dissociation constant of water with and has the value 10−14 mole2/l2. The pH is defined as −log [H+]; thus, pure water has a pH of 7. (These numbers are correct at 23 °C and slightly different at other temperatures.)
A base accepts (removes) hydrogen ions (H+) from the solution, or donates hydroxide ions (OH−) to the solution. Both actions will lower the concentration of hydrogen ions, and thus raise pH. By contrast, an acid donates H+ ions to the solution or accepts OH−, thus lowering pH.
The pH of a solution can be calculated. For example, if 1 mole of sodium hydroxide (40 g) is dissolved in 1 liter of water, the concentration of hydroxide ions becomes [OH−] = 1 mole/l. Therefore [H+] = 10−14 mol/l, and pH = −log 10−14 = 14.
Neutralization of acids
When dissolved in water, sodium hydroxide decomposes into hydroxide and sodium ions:
and similarly, hydrochloric acid forms hydronium and chloride ions:
When the two solutions are mixed, the H+ and OH− ions combine to form water molecules:
If equal amounts of NaOH and HCl (measured in moles, not grams) are dissolved, the base and the acid exactly neutralize, leaving only NaCl (table salt) in solution.
Alkalinity of non-hydroxides
Both sodium carbonate and ammonia are bases, although neither of these substances contains OH− groups. That is because both compounds accept H+ when dissolved in water:
Neutralization
Neutralization is a chemical reaction, also called a water forming reaction, in which an acid and a base react and produce salt and water. In other words, you can say that neutralization is the combinaton of hydrogen ions H+ and hydroxide ions OH- (or oxide ions O2-) to form water molecule H2O. In the process, a salt is formed. Neutralization is exothermic, meaning it produces heat.
Most generally, the following occurs:
Acid + Base --> Salt + Water : ΔH = −C < 0
As an example — the reaction between hydrochloric acid and sodium hydroxide solutions:
hydrochloric acid + sodium hydroxide --> sodium chloride + water
HCl(aq) + NaOH(aq) --> NaCl(aq) + H2O(l)
Since the HCl and NaOH dissociate into ions in solution, the ionic equation is:
H+ + Cl− + Na+ + OH− --> Na+ + Cl− + H2O(l)
And since the sodium and chloride ions are just spectator ions not involved in the reaction, the net equation becomes:
H+ + OH− --> H2O(l) : ΔHr = −56 kJ/mol
This illustrates why neutralization reactions are also referred to as water forming reactions. Of course the sodium and chloride ions are still in solution so the result is pH neutral salt water.
Chemical titration methods are used for analyzing acids or bases to determine the unknown concentration. A pH meter can be used to determine the point of neutralization or a pH indicator which shows the point of neutralization by a distinct color change can be used. Simple stoichiometric calculations with the known volume of the unknown and the known volume and molarity of the added chemical gives the molarity of the unknown.
Excess gastric acid in the stomach, acid indigestion, is typically neutralized by the ingestion of sodium bicarbonate (NaHCO3) or other neutralizing agent (antacids).
Phenolphthalein
phenophthalein structure
Phenolphthalein is a sensitive pH indicator with the formula C20H14O4. Often used in titrations, it turns from colorless in acidic solutions to pink in basic solutions, the color change occurring between pH 8 and pH 10. If the concentration of indicator is particularly strong, it can appear purple.
In strongly basic solutions, phenolphthalein's pink color undergoes a rather slow fading reaction and becomes colorless again. In other words, the molecule has three forms:
The fading reaction is sometimes used in undergraduate classes for the study of reaction kinetics.
Phenolphthalein is insoluble in water, and is usually dissolved in alcohols for use in experiments. It is itself a weak acid, which can lose H+ ions in solution. The phenolphthalein molecule is colorless. However, the phenolphthalein ion is pink. When a base is added to the phenolphthalein, the atom ↔ ions equilibrium shifts to the ionization because H+ ions are removed, (by Le Chatelier's principle).
Phenolphthalein has been used for over a century as a laxative, but is now being removed from the market because of concerns over carcinogenity. However, the small amounts usually used in experiments are harmless. Phenolphthalein is also commonly used in a mixture, primarily by forensic scientists, to test for the presence of blood.
Acid Base Theories
Common acid-base theories
Lavoisier's definition
The first scientific definition was proposed by the French chemist Antoine Lavoisier.
Since Lavoisier's knowledge of strong acids was mainly restricted to the oxyacids, which tend to contain central atoms in high oxidation states surrounded by oxygen, such as HNO3 and H2SO4, and since he was not aware of the true composition of the hydrohalic acids, HCl, HBr, and HI, he defined acids in terms of their containing oxygen, which in fact he named from Greek words meaning "acid-former". When the elements chlorine, bromine, and iodine were identified and the absence of oxygen in the hydrohalic acids was established by Sir Humphry Davy in 1810, this definition had to be rejected.
The Arrhenius definition
Svante Arrhenius provided the first modern definition of acids and bases in 1884. In water, a dissociation takes place:
H2O → H+ + OH-
A compound causing an increase in H+ and a decrease in OH- is an acid and one causing the reverse is a base.
An Arrhenius acid, when dissociated in water, typically yields positively-charged hydrogen ion and a complementary negative ion.
An Arrhenius base, when dissociated in water, typically yields negatively-charged hydroxide ion and a complementary positive ion.
The positive ion from a base can form a salt with the negative ion from an acid. For example, two moles of the base sodium hydroxide (NaOH) can combine with one mole of sulfuric acid (H2SO4) to form two moles of water and one mole of sodium sulfate.
2NaOH + H2SO4 → 2H2O + Na2SO4 .
Precaution Steps:
- To fill a buret, close the stopcock at the bottom and use a funnel. You may need to lift up on the funnel slightly, to allow the solution to flow in freely.
- Before titrating, condition the buret with titrant solution and check that the buret is flowing freely. To condition a piece of glassware, rinse it so that all surfaces are coated with solution, then drain. Conditioning two or three times will insure that the concentration of titrant is not changed by a stray drop of water.
- Check the tip of the buret for an air bubble. To remove an air bubble, whack the side of the buret tip while solution is flowing. If an air bubble is present during a titration, volume readings may be in error.
-
Place the tip of the pipette in the solution and release your grip on the bulb to pull solution into the pipette. Draw solution in above the mark on the neck of the pipette. If the volume of the pipette is larger than the volume of the pipette bulb, you may need to remove the bulb from the pipette and squeeze it and replace it on the pipette a second time, to fill the pipette volume completely.
-
After the solid is completely dissolved, very carefully fill the flask to the 250 mL mark. Move your eye to the level of the mark on the neck of the flask and line it up so that the circle around the neck looks like a line, not an ellipse. Then add distilled water a drop at a time until the bottom of the meniscus lines up exactly with the mark on the neck of the flask. Take care that no drops of liquid are in the neck of the flask above the mark.
Conclusion:
After the experiment, we found that the concentration of NaOH is 0.1M
References:
- http://en.wikipedia.org/wiki/Acid
- http://en.wikipedia.org/wiki/acid-base_reaction_theories
- http://en.wikipedia.org/wiki/Base
- http://en.wikipedia.org/wiki/Neutralization
- http://en.wikipedia.org/wiki/PH_indicator
- http://en.wikipedia.org/wiki/Base_(chemistry)
-
Fundamental of Analytical Chemistry 7th edition by Skoog, West and Holler, Saunders Publishers. 1996
-
Laboratory Experiments for Chemistry, A Basic Introduction, 4th edition, by Wynn and Joppich, Wadsworth Inc, 1987
Questions:
(1) Calculate the concentration of NaOH solution.
Concentration of base (diluted solution)
M1V1 = M2V2
M1(12.40) = (0.01)(25)
M = (0.01)(25)/12.40
= 0.0202M
The concentration of the diluted NaOH is 0.0202 molar, which is approximately 0.02 molar. From the molar of the diluted NaOH, we compare the concentration of the original NaOH,
Concentration of base (original solution)
M1V1 = M2V2
M1(50) = (0.0202)(250)
M = (0.0202)(250)/50
= 0.0101M
The concentration of the NaOH solution used is 0.0101 molar.
(2) Distinguish between acid strength and acid concentration.
Acid strength is the percentage of ionization of the acid when dissolve in water while acid concentration is the amount of dissolved acidic solutes in the solution.
(3) Distinguish between a weak base and an insoluble base.
A weak base is a chemical base that does not ionize fully in an aqueous solution. This results in a relatively low pH level. Weak bases exist in equillibrium much in the same way as weak acids do, with a Base Ionization Constant (Kb) indicating the strength of the base. Not many metal hydroxides are soluble; the ones that are comprise the strong soluble bases. Hydroxides that are only slightly soluble in water (such as calcium hydroxide or iron(III) hydroxide) are strong bases, because whatever amount does dissolve dissociates completely into the ions. So, we can say that weak base has a lower pH level compared to insoluble base because weak base does not ionize fully in aqueous solution, whereas insoluble base are strong base because most of them ionize fully in water.