BIOC 2200 – Experiment 1: Buffers and the effect on amino acids Aiden Forsyth 1009696399 Wednesday PM Bench 13 Results Questions A1. By the Henderson-Hasselbalch equation (Plumber, 2004). pH=pka+log〖([HA])/([A^-])〗 n_([HA])=M*V=0.10 mol/L*0.050L=0.0050 mol n_[A^- ] =n_NaOH=M*V=0.20 mol/L*0.001L=0.0002 mol Since acetic acid is transformed to sodium acetate there is a consumption of [HA] by the amount of moles of sodium hydroxide. pH=4.76+log〖0.0002/(0.0050-0.0002)〗=4.76-1.38=3.38 A3. To completely dissociate all of the acetic acid with sodium hydroxide to sodium acetate there must be zero moles of acetic acid remaining (Plumber, 2004). From A1 we know that the number of moles of acetic acid is While the pka for acetic acid can be determined in Graph 1 to be 4.73 the volume of sodium hydroxide in 3mL short of the calculated 13mL point. The published value of 4.76 is similar to the observed pka of acetic acid (Horton et al., 2006). The pka of boric acid was observed to be 9.27 at 14mL of sodium hydroxide added which is consistent with the published value of 9.27 (Silberberg, 2010). Boric acid would be a good buffer for an experiment conducted at a pH of 8.5 because the mixed solution would have a similar concentration of acid and conjugate base. Thereby resisting a change to its pka or 9.27. Acetic acid would have very little resistance to a change in the experiments pH because its pka is 4.76 thus being an already defeated buffer by the time we start the experiment at a pH of To fully dissociate an amino acid is to ionize all of the present ionizable groups (Horton et al., 2006). To such an end the environment’s pH must be higher than the highest pKa so as to deprotonate the last remaining group. The pKas for arginine are 1.8 for the alpha carboxyl group, 9.0 for the alpha amino group and 12.5 for the R group. So the aim is to have an environment with a pH higher than 12.5 to fully dissociate arginine. The corresponding volume of NaOH require to have a pH of 12.52 was 79mL. There must be three equivalents of NaOH to dissociate all three of the groups found in arginine where as in acetic and boric acid there are only two groups and therefore only need two NaOH equivalents. Valine is similar where it has two groups that can be ionized with pKas of 2.3 for the alpha carboxyl group and 9.7 for the alpha amino group (Plumber, 2004). For 1 mol of valine the required the volume of 0.20M NaOH needed is pH=pKa+log (〖[A〗^-])/[HA] 〖[A〗^-]=〖10〗^((9.8-9.7) )*1=1.3mol so 1.3 mol of NaOH is required D3. The two forms of threonine found in milk at a pH of 6.5 would exist in a ratio based on their isoelectric point (Hine et Martin, 2014). pI=(∑▒pKa)/2=(2.1+9.1)/2=5.6 Because the environment’s pH is less than the isoelectric point the two forms of threonine present are the zwitterion and cation forms. The ratio of the two being pH=〖pKa〗_cation+log (〖[A〗^-])/[HA]
As shown in Fig. 5, the final pH of the NaClO-NH3 solution after simultaneous removal are 5.4, 6.9, 7.2, 7.5, 8.5, 9.6, 10.7, 11.5 and 12.8 with respect to the initial pH of 5, 6, 7, 8, 9, 10, 11, 12 and 13, from which, an interesting law can be concluded as that if the initial pH is an acidic, the final pH is slightly increased; but if the initial pH is an alkaline, the final pH is declined. NaClO-NH3 is macromolecule compounds with a large inter surface area. It contains abundant functional groups such as hydroxyl (OH), carboxyl (COO), quinone, amino (–NH2), etc, which determines that NaClO-NH3 is a salt of strong base and weak acid, as well the ionization equilibrium and hydrolytic equilibrium would be complicated. When the pH of the NaClO-NH3 solution was acidic, the functional groups such as OH, COO and NH2- would react with H+ to generate the NH3 sediment, resulting in a decrease of inter surface area owing to the block and a great loss of NaClO-NH3, then the NOx removal as well as the duration time was decreased. As for the increase of the final pH in the acidic conditions, this was a result of the consumption of H+ by NaClO. The decrease of the
The citric acid cycle is an amphibolic pathway. It utilises both anabolic and catabolic reactions; the first reaction of the cycle, in which oxaloacetate (a four carbon compound) condenses with acetate (a two carbon compound) to form citrate (a six carbon compound) is typically anabolic. The production of the isomeric isocitrate is simply intramolecular rearrangement. The subsequent two reactions are typically catabolic, producing succinate (a four carbon compound), which is then oxidised, forming fumarate (a four carbon compound). Water addition produces malate and then oxidised for regeneration of oxaloacetate. Thus the cycle can be seen to exhibit both anabolic and catabolic processes to form its intermediates.
The sought to determine the effects of mixing various levels of acids and bases to see which combination would have the most explosive reaction, and measure the resulting pH levels. I did this by testing an assortment of different pH levels of acids and bases, mixing them together and measuring the results. Most of the experiments resulted in a pH neutral solution, except for the Sulfuric Acid and the Sodium Hydroxide. By far, the Sulfuric Acid was the most explosive, followed by the Citric and Acetic acid.
Then the OH- and H+ molecules combine to form H2O water molecules and a sodium benzoate compound. Sodium benzoate then becomes insoluble in diethyl ether and soluble water. However, naphthalene does not react with the NaOH because it is not soluble in NaOH. Naphthalene and sodium benzoate differ in solubility and can therefore be separated into an aqueous layer and an organic layer. The diethyl ether is nonpolar and naphthalene has a low polarity so they are soluble together. The sodium benzoate is soluble in water. Diethyl ether is less dense than water so it stays on the top organic layer with naphthalene and sodium benzoate and water separate to the bottom aqueous layer of the separatory
ranging from 50 cm³ of acid and no water, to 12.5 cm³ of acid and 37.5
an unknown amino acid. A titration curve is the plot of the pH versus the volume
It is important however to note that the NH4 and K ions are still in
= = pH 1 2 3 Average Rate of Reaction (cm3/s): 0 - 0. 3 0 0 0 0 0.000 5 0 0 0 0
Buffers are comprised of weak acids and their conjugate bases. In the food industry, buffers are commonly used to protect changes in pH of food stuffs. The two conjugate components of the buffer resist changes in pH by absorbing the addition of any hydrogen or hydroxyl ions (Christen and Smith 2000). When weak acids and their conjugate bases are at equilibrium, their concentrations can be expressed in terms of the dissociation constant, Ka. For a strong acid, the Ka value is greater than 1 x 10-2M and less than 1 x 10-2M for a weak acid (Thompson and Dinh 2009). For weak acids, the relationship of hydrogen ion concentration and pH can be defined by the Henderson-Hasselbach equation (Thompson and Dinh 2009).
Acid-Base balance is the state of equilibrium between proton donors and proton acceptors in the buffering system of the blood that is maintained at approximately pH 7.35 to 7.45 under normal conditions in arterial blood. It is important to regulate chemical balance or homeostasis of body fluids. Acidity or alkalinity has to be regulated. An acid is a substance that lets out hydrogen ions in solution. Strong acid like hydrochloric acid release all or nearly all their hydrogen ions and weak acids like carbonic acid release some hydrogen ions.
To determine the pH titration curve for both strong acid – base titration and weak – acid base titration.
Study on the minimum inhibitory concentration of acetic acid on E. coli and S. aureus
The titration of a weak acid with a strong base produces a titration curve as above.
The simplest experiment for this type of situation would be to use red and blue litmus paper to distinguish between acids, bases and salts. Hydrochloric acid (HCl) makes blue litmus paper change color going from blue to red, making it an acid. Sodium hydroxide (NaOH) makes red litmus paper change color going from red to blue, making it a base. Sodium chloride solution (NaCl) is neutral, since it would only soak blue and red litmus paper, considering that it is a by product of when an acid and a base mix together, neutralizing each other.
The purpose of this experiment is to use our knowledge from previous experiments to determine the exact concentration of a 0.1M sodium hydroxide solution by titration (Lab Guide pg.141).