Ionic Equilibria - Lab Report Example

Lab report: Ionic equilibria (Word count: 1456) Laboratory report: Ionic equilibria Aim of the experiment. To carry out an acid-base pH titration, plot a suitable graph and determine the concentration of the ethanoic acid solution. Introduction. According to Brønsted-Lowry theory an acid can be defined as a "proton donor" while a base can be called a "proton acceptor." Acids are in most cases divided into groups such as "strong" and "weak." One measure of the most commonly used measure of acid strength is acid-dissociation equilibrium constant, Ka, for that particular acid (Atkins and Paula 2001, Pg. 45). A strong acid has a big Ka value. For example hydrochloric acid HCl: Ka = 1 x 103 While a weak acid has a small Ka value. For example ethanoic acid (Harris 2006, Pg. 23). CH3CO2H: Ka = 1.8 x 10-5 However, when it very small, then we talk about a weak acid. H2O: Ka = 1.8 x 10-16 Reaction between acid and bases yield two products: water, and an ionic compound known as a salt. This type of reaction is called a neutralization reaction. Strong and weak acids Generally acid dissolves in water as shown in the equation below, However, strong acid react only in the forward direction for example HCl as shown in the equation below. While, weak acids are the acids which do not fully ionize in water, For example, ethanoic acid as shown below (Atkins and Paula 2001, Pg. 46). . In order to measure the dissociation constant of an acid a dissociation constant is always used. For example, when using ethanoic acid. The Ka will be given as. Strong and Weak bases A strong base is one which dissolves completely in water for example NaOH. A weak base does not fully ionize in water for example ammonia as shown below, In order to compare the strength of weak bases then equilibrium constant Kb is used as shown below (Atkins and Paula 2001, Pg. 47). , An acid dissolves in water, a proton (hydrogen ion) is transferred to a water molecule to produce a hydroxonium ion and a negative ion depending on what acid you are starting from (Garland 2006, Pg. 31). TITRATION Since acids and bases readily react with each other, it is experimentally quite easy to find the amount of acid in a solution by determining how many moles of base are required to neutralize it. This operation is called titration, and you should already be familiar with it from your work in the Laboratory (Harris 2006, Pg. 06). We can titrate an acid with a base, or a base with an acid. The substance whose concentration we are determining is the substance being titrated; the substance we are adding in measured amounts is the titrant. The idea is to add titrant until the solution has been exactly neutralized; at this point, the number of moles of titrant added tells us the concentration of base (or acid) in the solution being titrated (Harris 2006, Pg. 8). These definitions carry a very important implication: a substance cannot act as an acid without the presence of a base to accept the proton, and vice versa. A reaction of an acid with a base is thus a proton exchange reaction; if the acid is denoted by AH and the base by B, then we can write a generalized acid-base reaction as AH + B − → A− + BH+ But the product BH+ is now capable of losing its newly-acquired proton to another acceptor, and is therefore potentially another acid: acid1 + base2 − → base1 + acid2 In this schematic reaction, base1 is conjugate to acid1, and acid2 is conjugate to base2. The term conjugate means “connected with”, the implication being that any species and its conjugate species are related by the gain or loss of one proton. THE HYDRONIUM ION The Arrhenius view of an acid is a substance that dissociates in water to produce a hydrogen ion. There is a serious problem with this, however: the hydrogen ion is no more than a proton, a bare nucleus. Although it carries only a single unit of positive charge, this charge is concentrated into a volume of space that is only about a hundred-millionth as large as the volume occupied by the smallest atom. Owing to its extremely small size, the proton will be attracted to any part of a nearby atom or molecule in which there is an excess of negative charge. Such places exist on any atom that possesses non-bonding electrons, and here that protons attach themselves to the acceptor atom by forming a shared-electron (coordinate) bond with the lone pair (Silbey 2006, Pg. 42). Interestingly, experiments indicate that the proton does not stick to a single H2 O molecule, but changes partners many times per second. This molecular promiscuity, a consequence of the uniquely small size and mass the proton, allows it to move through the solution by rapidly hopping from one H2 O molecule to the next, creating a new H3 O+ ion as it goes. The overall effect is the same as if the H3 O+ ion itself were moving. Similarly, a hydroxide ion, which can be considered to be a “proton hole” in the water, serves as a landing point for a proton from another H2 O molecule, so that the OH− ion hops about in the same way (Atkins and Paula 2001,Pg. 50). Because hydronium- and hydroxide ions can “move without actually moving” and thus without having to plow their way through the solution by shoving aside water molecules as do other ions, solutions which are acidic or alkaline have extraordinarily high electrical conductivities (Atkins and Paula 2001,Pg. 55). METHODOLOGY Apparatus Retort stand and clamp Wash bottle 2 beaker 250cm3 3 beaker 100cm3 pH electrode pH meter Burette 20.0cm3 bulb pipette and filler Chemicals Ethanoic acid 50cm3 1.00M sodium hydroxide 100 cm3 Buffer pH 4.00 25cm3 Buffer pH 7.00 25cm3 Procedure Once the apparatus was set-up, the pH electrode and meter were properly calibrated. The pH electrode was placed in the alkali in a position where the acid would not be titrated directly on top of the electrode. The acid was then added in small increments and change in pH was recorded as observed in the table of results. The electrode was also used to stir the mixture before a reading was noted in the table. Once the full 50cm3 of acid was added and the final reading taken the apparatus was cleaned up and a graph tabulated to find out the molarity of the acid used. Calibration of pH meter with buffers A buffer solution is a type of solution that resists change even when a small quantities of acid or base is added to it (Silbey 2004, Pg. 35). In order to calibrate a pH meter a 4.01 and 7.00 pH buffer solution is needed. The electrode of the pH meter were rinsed with deionized water to remove traces of all storage solution. They were then immersed in 7.00 and 4.01 pH buffer solution. The readings were then adjusted accordingly (Silbey 2004, Pg. 34). Chemical reaction CH3COOH + NaOH ==> CH3COONa + H2O The titration curve This a reaction between a strong base (sodium hydroxide) and weak acid (Ethanoic acid). The expected graph for this kind of reaction is as shown below. The theoretical end point of the reaction pH + pOH = 14. The pOH of NaOH is given by pOH = - log [1] =0 Therefore the pH of ethanoic acid at equivalence point is 14. RESULTS Analysis of the graph 1. The original pH is higher before the addition of the weak acid. 2. There is a sharp rise in pH at the start of the titration. This is can be attributed to the fact that the anion of weak acid becomes the common ion hence reduces ionisation of the acid. 3. Then after the sharp increase at the beginning there is a slow rise because the solution is now acting as a buffer hence it resists any changes in pH. The trend continues until the base overcomes the buffer. After the sharp increase at the beginning of the titration the curve only changes gradually. 4. At the equivalence point the pH is greater than 7, this is because all acid (HA) has been converted to a conjugate base (-A) by the continuous addition of NaOH and hence the equilibrium shifts backwards and produces more hydroxide. A−+H2O⇌AH+OH− Determination of the concentration of the acid. From the graph the equivalent point is 6.5. But pH = - log [H3O+]. 6.5 = - log [CH3CHOH] [CH3CHOH]= 3.16* 10 -7 Percentage error The theoretical concentration is 14 = - log [CH3CHOH] [CH3CHOH] = 1*10-14 Percentage error = (1*10-14 )/ (3.16* 10 -7) = 3.16 * 10 -8 References Atkins, P. Paula, J. (2001). Physical Chemistry. New York: Freeman. Harris, D. C. (2006). Qualitative chemical analysis. New York: Freeman Silbey, R. J. (2004). Physical chemistry. Chicago: Wiley. Garland, C. (2008). Experiments in Physical Chemistry. New York: Wiley and Sons. Read More