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Home , Bromic acid, Fluorosulfuric acid, Hydrobromic acid

AND BASES Properties of acids 1. Acids have a sour taste. 2. Acids are corrosive. 3. Acids change the color of certain vegetable dyes, such as litmus, from blue to red. 4. Acids lose their acidity when they are combined with alkalies. The name "" comes from the Latin acidus, which means "sour," and refers to the sharp odor and sour taste of many acids. Examples: Vinegar tastes sour because it is a dilute solution of in water. Lemon juice tastes sour because it contains citric acid. Milk turns sour when it spoils because lactic acid is formed, and the unpleasant, sour odor of rotten meat or butter can be attributed to compounds such as butyric acid that form when fa`t spoils. Acids, Bases, and pH Calculations In chemistry, pH is a measure of the acidity or basicity of an aqueous solution. Solutions with a pH less than 7 are said to be acidic and solutions with a pH greater than 7 are basic or alkaline. Pure water has a pH very close to 7. There are several ways to define acids and bases, but pH only refers to ion concentration and is only meaningful when applied to aqueous (water-based) solutions. When water dissociates it yields a hydrogen ion and a hydroxide. + - H2O ↔ H + OH When calculating pH, remember that [] refers to molarity, M. + - -14 Kw = [H ][OH ] = 1x10 at 25°C for pure water [H+] = [OH-] = 1x10-7 Acidic Solution: [H+] > 1x10-7 Basic Solution: [H+] < 1x10-7 Calculate pH and [H+] + pH = - log10[H ] [H+] = 10-pH Example: Calculate the pH for a specific [H+].C4alculate pH given [H+]= 1.4 x 10-5 M + pH =- log10[H ] -5 pH = log10(1.4 x 10 ) pH = 4.85 Example: Calculate [H+] from a known pH. Find [H+] if pH = 8.5 [H+] = 10-pH [H+] = 10-8.5 [H+] = 3.2 x 10-9 M What is the pH of a solution with [H+] = 1 x 10-6 M. Solution pH is calculated by the formula pH = - log [H+] Substitute [H+] with the concentration in the question. pH = - log (1 x 10-6) pH = -(-6) pH = 6 Answer The pH of the solution is 6. What is the pH of a 0.025 M solution of Hydrobromic Acid? Solution Hydrobromic Acid or HBr, is a strong acid and will dissociate completely in water to H+ and Br-. For every mole of HBr, there will be 1 mole of H+, so the concentration of H+ will be the same as the concentration of HBr. Therefore, [H+] = 0.025 M. pH is calculated by the formula pH = - log [H+] Enter the concentration found before pH = - log (0.025) pH = -(-1.602) pH = 1.602 Answer The pH of a 0.025 M solution of Hydrobromic Acid is 1.602.

Calculating pH of a strong base. Question What is the pH of a 0.05 M solution of Potassium Hydroxide? Solution Potassium Hydroxide or KOH, is a strong base and will dissociate completely in water to K+ and OH-. For every mole of KOH, there will be 1 mole of OH-, so the concentration of OH- will be the same as the concentration of KOH. Therefore, [OH-] = 0.05 M. Since the concentration of OH- is known, the pOH value is more useful. pOH is calculated by the formula pOH = - log [OH-] Enter the concentration found before pOH = - log (0.05) pOH = -(-1.3) pOH = 1.3 The value for pH is needed and the relationship between pH and pOH is given by pH + pOH = 14 pH = 14 - pOH pH = 14 - 1.3 pH = 12.7 Answer The pH of a 0.05 M solution of Potassium Hydroxide is 12.7. Exercises CALCULATING THE pH AND pOH OF STRONG ACIDS AND BASES • This is relatively easy because the species have completely dissociated • Only needs to know the original concentration of the acid or base Example 1 Calculate the pH of 0.1M . HCl (a strong monoprotic acid) is fully dissociated. HCl ——> H+ (aq) + Cl¯(aq) The [ H+] is therefore the same as the original concentration of HCl i.e. 0.1M. pH = - log10 [H+] = - log10 (10-1) = 1 ANS. 1 Example 2 Calculate the pH of 0.001M sodium hydroxide. Sodium hydroxide (a strong base) is fully dissociated. Na+OH¯ ——> Na+ (aq) + OH¯(aq) The [OH¯] is therefore the same as the original concentration of NaOH i.e. 0.001M. pOH = - log10 [OH¯] = - log10 (10-3) = 3 and pH = 14 - pOH = 14 - 3 = 11 ANS. 11 The dissociation constant for a weak acid (Ka) A weak monobasic acid (HA) dissociates in) H+ Acid-base A4 5 Q.4 Calculate the pH and pOH of the following solutions. a) HCl; 0.1M, 0.5M b) H2SO4; 0.1M, 0.5M c) KOH; 0.1M d) NaOH; 2M, 0.0005M e) The solution remaining when 30 cm3 of 0.100M NaOH has been added to 20 cm3 of 0.200M HCl f) The solution remaining when 24.9 cm3 of 0.100M NaOH has been added to 25 cm3 of 0.100M HCl

Acid strength The strength of an acid refers to its ability or tendency to lose a proton (H+). A strong acid is one that completely ionizes (dissociates) in a solution. In water, one mole of a strong acid HA dissolves yielding one mole + + of H (as hydronium ion H3O ) and one mole of the conjugate base, A−. Essentially none of the non-ionized acid HA remains. Examples of strong acids are hydrochloric acid (HCl), (HI), hydrobromic acid (HBr), perchloric acid (HClO4), nitric acid (HNO3) and (H2SO4). In aqueous solution each of these essentially ionizes 100%. In contrast, a weak acid only partially dissociates. Examples in water include carbonic acid (H2CO3) and acetic acid (CH3COOH). At equilibrium both the acid and the conjugate base are present in solution. Stronger acids have a larger acid dissociation constant (Ka) and a smaller logarithmic constant (pKa = - log Ka) than weaker acids. The stronger an acid is, the more easily it loses a proton, H+. Two key factors that contribute to the ease of deprotonation are the polarity of the H—A bond and the size of atom A, which determines the strength of the H—A bond. Acid strengths also depend on the stability of the conjugate base.

While Ka measures the strength of an acidic molecule, the strength of an aqueous acid solution is measured by pH, which is a function of the concentration of hydronium ion in solution. The pH of a simple solution of an acid in water is determined by both Ka and the acid concentration. For weak acid solutions it depends on the degree of dissociation, which may be determined by an equilibrium calculation. For concentrated solution of strong acids with pH less than about zero, the is a better measure of acidity than the pH. Sulfonic acids, which are organic , are a class of strong acids. A common example is p-toluenesulfonic acid (tosylic acid). Unlike sulfuric acid itself, sulfonic acids can be solids. In fact, polystyrene functionalized into polystyrene sulfonate is a solid strongly acidic plastic that is filterable. are acid solutions which are more acidic than 100% sulfuric acid.[1] Examples of superacids are fluoroantimonic acid, and perchloric acid. Superacids can permanently protonate water to give ionic, crystalline hydronium "salts". They can also quantitatively stabilize carbocations. Strong acids A strong acid is an acid that ionizes completely in an aqueous solution by losing one proton, according to the equation HA(aq) → H+(aq) + A−(aq) For sulfuric acid which is diprotic, the "strong acid" designation refers only to dissociation of the first proton + H2SO4(aq) → H (aq) + − HSO4 (aq) More precisely, the acid must be stronger in aqueous solution than hydronium ion, so strong acids are acids with a pKa < −1.74. An example is HCl for [2] which pKa = -6.3. This generally means that in aqueous solution at standard temperature and pressure, the concentration of hydronium ions is equal to the concentration of strong acid introduced to the solution. In all other acid-water reactions, dissociation is not complete, so will be represented as an equilibrium, not a completed reaction. The typical definition of a weak acid is any acid that does not dissociate completely. The difference separating the acid dissociation constants of strong acids from all other acids is so small that this is a reasonable demarcation. Due to the complete dissociation of strong acids in aqueous solution, the concentration of hydronium ions in the water is equal to the total concentration (ionized and un-ionized) of the acid introduced to solution: [H+] = − [A ] = [HA]total and pH = −log[H+]. Determining The strength of an acid, in comparison to other acids, can be determined without the use of pH calculations by observing the following characteristics: 1. Electronegativity: The higher the electronegativity of a conjugate base in the same period, the more acidic. In other words, the more electronegative A- is, more acidic (where HA → H+ + A−). 2. Atomic Radius: With increasing atomic radius, acidity also increases. For example, HCl and HI, both strong acids, ionize 100% in water to become their respective ionic constituents. However, HI is stronger than HCl. This is because the atomic radius of an atom of is much larger than that of a chlorine atom. As a result, the negative charge over the I− anion is dispersed over a larger electron cloud and its attraction for the proton (H+) is not as strong as the same attraction in HCl. Therefore, HI is ionized (deprotonated) more readily. 3. Charge: The more positively charged a species is, the more acidic (neutral molecules can be stripped of protons more easily than anions, and cations are more acidic than comparable molecules). 4. Equilibrium: The strength of an acid can also be defined by the equilibrium position of its dissociation reaction: + HA(aq)+ H2O(l) → H3O (aq) + A−(aq) In a strong acid, equilibrium lies far to the right, meaning that almost all of the original HA is dissociated at equilibrium. A strong acid yields a weak conjugate base (A−), so a strong acid is also described as an acid whose conjugate base is a much weaker base than water.[3] Common strong acids This is a list of strong acids with pKa < -1.74, which is stronger than hydronium ion, from strongest to weakest.

 Perchloric acid HClO4 (pKa ≈ −10)[4]

 Hydroiodic acid HI (pKa = −9.3)[2]

 Hydrobromic acid HBr (pKa = −8.7)[2]

 Hydrochloric acid HCl (pKa = −6.3)[2]

 Sulfuric acid H2SO4 (first dissociation only, pKa1 ≈ −3)[5]  p-Toluenesulfonic acid (pKa = −2.8) Organic soluble strong acid

 Methanesulfonic acid (pKa = −1.92) Liquid organic strong acid[6] Almost strong acids These do not meet the strict criterion of being more acidic + than H3O , although in very dilute solution they dissociate almost completely, so sometimes they are included as "strong acids" +  Hydronium ion H3O (pKa = -1.74). Hydronium is often used as an approximation of the state of protons in water.

 Nitric acid HNO3 (pKa = - 1.64)[5]

HClO3 (pKa = - 1.0)[5]

 Some chemists include [7] bromic acid (HBrO3), [7] (HBrO4), [7] (HIO3), and [7] periodic acid (HIO4) as strong acids, although these are not universally accepted as such. Extremely strong acids (as protonators) (Strongest to weakest)

 Fluoroantimonic acid H[SbF6]

 Magic acid FSO3HSbF5

 Carborane H(CHB11Cl11)

 Fluorosulfuric acid FSO3H [6] (pKa = -6.4)

CF3SO3H (pKa = -5.9)[6] Weak acids Most acids are weak acids. A weak acid is an acid that dissociates incompletely, releasing only some of its hydrogen atoms into the solution. Thus, it is less capable than a strong acid of donating protons. These acids have higher pKa than strong acids, which release all of their hydrogen atoms when dissolved in water. Examples of weak acids include acetic acid (CH3COOH) and oxalic acid (H2C2O4). Dissociation Weak acids ionize in water solution to only a moderate extent; that is, if the acid was represented by the general formula HA, then in aqueous solution a significant amount of undissociated HA still remains. Weak acids in water dissociate as:

The strength of a weak acid is represented as either an equilibrium constant or as a percent dissociation. The equilibrium concentrations of reactants and products are related by the acid dissociation constant expression, (Ka):

The greater the value of Ka, the more the formation of H+ is favored, and the lower the pH of the solution. The Ka of weak acids varies between 1.8×10−16 and 55.5. Acids with a Ka less than 1.8×10−16 are weaker acids than water. The other way to measure acid strength is to look at its fractional dissociation, which is symbolized as α (alpha) and which can range from 0% < α < 100%. The dissociation ratio is defined as

Unlike Ka, α is not constant and does depend on the [HA]. In general α will increase as [HA] decreases. Thus acids become stronger as they are diluted. If acids are polyprotic, then each proton will have a Ka. For example: H2CO3 + H2O → − + HCO3 + H3O has two Ka values because it has two acidic protons. The first Ka value is −7 4.46×10 (pKa1 = 6.351) and −11 the second is 4.69×10 (pKa2 = 10.329). Calculating the pH of a weak acid solution The pH of a solution of a weak acid depends on the strength of the acid and the other components in the solution. In the simplest case, the weak acid is the only compound in water. In this case, the pH can be found from the concentration of the acid (symbolized as ), from the Acid dissociation constant (symbolized as ), and by solving for concentration of H+ (symbolized by x and represented more accurately as + H3O ). Below is a table that organizes the information. On the first line, the reaction is written. On the second line, the initial conditions are written below each compound. Note that a value of water is not given because its term (activity) in the expression is technically equal to 1, but is often (conveniently) omitted. The third line shows how the value changes as the reaction goes to equilibrium. Then the last line gives the equilibrium concentrations and is simply the sum of each column. HA H A− H O+ ( + 2 → ( + 3 aq) O(l) aq) (aq) Initial F — 0 0

Change -x — +x +x

Equilibr F - — x X ium x Applying the equilibrium line to the expression yields rearranging yields , which can be solved for x using the quadratic equation. The pH is then calculated as . Simplification However, if F is more than 1000× greater than Ka, then (1) the acid will not deprotonate much, (2) the value of x will be small, and therefore (3) F - x ≈ F. This simplifies the Ka expression to...

Solving for x yields

Then the pH = -log[H+]. The following equation then follows, but is only true if F >>> Ka

Comparison of the full and simplified methods

A certain weak acid has a Ka = 1×10−5 and the pH of two solutions needs to be found. One solution has a concentration of 0.10M and another has a concentration of 5×10−4M. The pH for both solutions will be calculated using both methods to yield 4 values, which will be compared. 0.1M Solution The full method gives the following quadratic: which gives x = 9.95×10−4 M and a pH = 3.00. The simplified method gives

So both methods yield the same result, but again F is more than 1000× greater than Ka. The next case does not have this condition and the results will differ. 5×10−4M Solution The full method gives the following quadratic: which gives x = 6.6×10−5 M and a pH = 4.18. The simplified method gives

Here, the results differ by 0.03 pH units. As F becomes closer in value to the Ka, then the difference will increase even more. However, in practice, it is rare to work with such dilute acids and the pH is also dependent on ionic strength and temperature. So in reality, the simplified method works well. /base pair It is often stated that "the conjugate of a weak acid is a strong base". This statement can be misleading. Most weak acids that textbooks discuss have weak (not strong) conjugate bases. Truly, only the very weakest of acids have strong conjugate bases. For example, if −5 a weak acid has a Ka = 10 , then its conjugate base would −9 have a Kb = 10 (from the −14 relationship Ka × Kb = 10 ), which certainly is not a strong base. A very weak acid with a −20 Ka = 10 would indeed have a strong conjugate base. Factors determining acid strength Polarity and the inductive effect Polarity refers to the distribution of electrons in a bond, the region of space between two atomic nuclei where a pair of electrons is shared. When two atoms have roughly the same electronegativity (ability to attract electrons) the electrons are shared evenly and spend equal time on either end of the bond. When there is a significant difference in electronegativities of two bonded atoms, the electrons spend more time near the nucleus of the more electronegative element and an electrical dipole, or separation of charges, occurs, such that there is a partial negative charge localized on the electronegative element and a partial positive charge on the electropositive element. Hydrogen is an electropositive element and accumulates a slightly positive charge when it is bonded to an electronegative element such as oxygen or . As the electron density on hydrogen decreases it is more easily abstracted, in other words, it is more acidic. Moving from left to right across a row on the periodic table elements become more electronegative (excluding the noble gases), and the strength of the binary acid formed by the element increases accordingly: [8] Formula Name pKa hydrofluoric HF 3.17

acid

H2O Water 15.7

NH3 ammonia 38

CH4 methane 48 The electronegative element need not be directly bonded to the acidic hydrogen to increase its acidity. An electronegative atom can pull electron density out of an acidic bond through the inductive effect. The electron-withdrawing ability diminishes quickly as the electronegative atom moves away from the acidic bond. The effect is illustrated by the following series of halogenated butanoic acids. Chlorine is more electronegative than bromine and therefore has a stronger effect. The hydrogen atom bonded to the oxygen is the acidic hydrogen. Butanoic acid is a carboxylic acid. [9] Structure Name pKa butanoic acid ≈4.8 or butyric acid 4- chlorobutanoic 4.5 acid 3- chlorobutanoic ≈4.0 acid 2- bromobutanoic 2.93 acid 2- 2.86 chlorobutanoic acid As the chlorine atom moves further away from the acidic O—H bond, its effect diminishes. When the chlorine atom is just one carbon removed from the carboxylic acid group the acidity of the compound increases significantly, compared to butanoic acid (a.k.a. butyric acid). However, when the chlorine atom is separated by several bonds the effect is much smaller. Bromine is much more electronegative than either carbon or hydrogen, but not as electronegative as chlorine, so the pKa of 2-bromobutanoic acid is slightly greater than the pKa of 2-chlorobutanoic acid.

Perchloric acid (HClO4) is an oxoacid and a strong acid. The number of electronegative atoms adjacent an acidic bond also affects acid strength. Oxoacids have the general formula HOX where X can be any atom and may or may not share bonds to other atoms. Increasing the number of electronegative atoms or groups on atom X decreases the electron density in the acidic bond, making the loss of the proton easier. Perchloric acid is a very strong acid (pKa ≈ -8) and completely dissociates in water. Its is HClO4 and it comprises a central chlorine atom with three chlorine-oxygen double bonds (Cl=O) and one chlorine- oxygen single bond (Cl—O). The singly bonded oxygen bears an extremely acidic hydrogen atom which is easily abstracted. In contrast, chloric acid (HClO3) is a weaker acid, though still quite strong (pKa = - 1.0), while chlorous acid (HClO2, pKa = +2.0) and (HClO, pKa = +7.53) acids are weak acids.[10] Carboxylic acids are organic acids that contain an acidic hydroxyl group and a carbonyl (C=O bond). Carboxylic acids can be reduced to the corresponding ; the replacement of an electronegative oxygen atom with two electropositive yields a product which is essentially non-acidic. The reduction of acetic acid to ethanol using LiAlH4 (lithium aluminium hydride or LAH) and ether is an example of such a reaction.

The pKa for ethanol is 16, compared to 4.76 for acetic acid.[9][11] Atomic radius and bond strength Another factor that contributes to the ability of an acid to lose a proton is the strength of the bond between the acidic hydrogen and the atom that bears it. This, in turn, is dependent on the size of the atoms sharing the bond. For an acid HA, as the size of atom A increases, the strength of the bond decreases, meaning that it is more easily broken, and the strength of the acid increases. Bond strength is a measure of how much energy it takes to break a bond. In other words, it takes less energy to break the bond as atom A grows larger, and the proton is more easily removed by a base. This partially explains why is considered a weak acid while the other hydrohalic acids (HCl, HBr, HI) are strong acids. Although fluorine is more electronegative than the other halogens, its atomic radius is also much smaller, so it shares a stronger bond with hydrogen. Moving down a column on the periodic table atoms become less electronegative but also significantly larger, and the size of the atom tends to dominate its acidity when sharing a bond to hydrogen. Hydrogen sulfide, H2S, is a stronger acid than water, even though oxygen is more electronegative than . Just as with the halogens, this is because sulfur is larger than oxygen and the H—S bond is more easily broken than the H—O bond. Corrosivity While strong acids are generally assumed to be the most corrosive, this is not always true. The carborane superacid H(CHB11Cl11), which is one million times stronger than sulfuric acid,[12][13] is entirely non-corrosive, whereas the weak acid hydrofluoric acid (HF) is corrosive and can dissolve, among other things, glass[14] and most metals. pH (TITRATION) CURVES

This page describes how pH changes during various acid- base titrations.

The equivalence point of a titration Sorting out some confusing terms When you carry out a simple acid-base titration, you use an indicator to tell you when you have the acid and alkali mixed in exactly the right proportions to "neutralise" each other. When the indicator changes colour, this is often described as the end point of the titration. In an ideal world, the colour change would happen when you mix the two solutions together in exactly equation proportions. That particular mixture is known as the equivalence point. For example, if you were titrating sodium hydroxide solution with hydrochloric acid, both with a concentration of 1 mol dm-3, 25 cm3 of sodium hydroxide solution would need exactly the same volume of the acid - because they react 1 : 1 according to the equation.

In this particular instance, this would also be the neutral point of the titration, because sodium chloride solution has a pH of 7. But that isn't necessarily true of all the salts you might get formed. For example, if you titrate ammonia solution with hydrochloric acid, you would get ammonium chloride formed. The ammonium ion is slightly acidic, and so pure ammonium chloride has a slightly acidic pH. That means that at the equivalence point (where you had mixed the solutions in the correct proportions according to the equation), the solution wouldn't actually be neutral. To use the term "neutral point" in this context would be misleading. Similarly, if you titrate sodium hydroxide solution with ethanoic acid, at the equivalence point the pure sodium ethanoate formed has a slightly alkaline pH because the ethanoate ion is slightly basic. To summarise:

 The term "neutral point" is best avoided.

 The term "equivalence point" means that the solutions have been mixed in exactly the right proportions according to the equation.

 The term "end point" is where the indicator changes colour. As you will see on the page about indicators, that isn't necessarily exactly the same as the equivalence point.

Simple pH curves All the following titration curves are based on both acid and alkali having a concentration of 1 mol dm-3. In each case, you start with 25 cm3 of one of the solutions in the flask, and the other one in a burette. Although you normally run the acid from a burette into the alkali in a flask, you may need to know about the titration curve for adding it the other way around as well. Alternative versions of the curves have been described in most cases. Titration curves for strong acid v strong base We'll take hydrochloric acid and sodium hydroxide as typical of a strong acid and a strong base.

Running acid into the alkali

You can see that the pH only falls a very small amount until quite near the equivalence point. Then there is a really steep plunge. If you calculate the values, the pH falls all the way from 11.3 when you have added 24.9 cm3 to 2.7 when you have added 25.1 cm3. Running alkali into the acid This is very similar to the previous curve except, of course, that the pH starts off low and increases as you add more sodium hydroxide solution.

Again, the pH doesn't change very much until you get close to the equivalence point. Then it surges upwards very steeply.

Titration curves for strong acid v weak base This time we are going to use hydrochloric acid as the strong acid and ammonia solution as the weak base.

Running acid into the alkali

Because you have got a weak base, the beginning of the curve is obviously going to be different. However, once you have got an excess of acid, the curve is essentially the same as before. At the very beginning of the curve, the pH starts by falling quite quickly as the acid is added, but the curve very soon gets less steep. This is because a buffer solution is being set up - composed of the excess ammonia and the ammonium chloride being formed. Notice that the equivalence point is now somewhat acidic ( a bit less than pH 5), because pure ammonium chloride isn't neutral. However, the equivalence point still falls on the steepest bit of the curve. That will turn out to be important in choosing a suitable indicator for the titration. Running alkali into the acid At the beginning of this titration, you have an excess of hydrochloric acid. The shape of the curve will be the same as when you had an excess of acid at the start of a titration running sodium hydroxide solution into the acid. It is only after the equivalence point that things become different. A buffer solution is formed containing excess ammonia and ammonium chloride. This resists any large increase in pH - not that you would expect a very large increase anyway, because ammonia is only a weak base.

Titration curves for weak acid v strong base We'll take ethanoic acid and sodium hydroxide as typical of a weak acid and a strong base.

Running acid into the alkali For the first part of the graph, you have an excess of sodium hydroxide. The curve will be exactly the same as when you add hydrochloric acid to sodium hydroxide. Once the acid is in excess, there will be a difference.

Past the equivalence point you have a buffer solution containing sodium ethanoate and ethanoic acid. This resists any large fall in pH. Running alkali into the acid

The start of the graph shows a relatively rapid rise in pH but this slows down as a buffer solution containing ethanoic acid and sodium ethanoate is produced. Beyond the equivalence point (when the sodium hydroxide is in excess) the curve is just the same as that end of the HCl - NaOH graph.

Titration curves for weak acid v weak base The common example of this would be ethanoic acid and ammonia.

It so happens that these two are both about equally weak - in that case, the equivalence point is approximately pH 7. Running acid into the alkali This is really just a combination of graphs you have already seen. Up to the equivalence point it is similar to the ammonia - HCl case. After the equivalence point it is like the end of the ethanoic acid - NaOH curve.

Notice that there isn't any steep bit on this graph. Instead, there is just what is known as a "point of inflexion". That lack of a steep bit means that it is difficult to do a titration of a weak acid against a weak base. A summary of the important curves The way you normally carry out a titration involves adding the acid to the alkali. Here are reduced versions of the graphs described above so that you can see them all together.

More complicated titration curves Adding hydrochloric acid to sodium carbonate solution The overall equation for the reaction between sodium carbonate solution and dilute hydrochloric acid is:

If you had the two solutions of the same concentration, you would have to use twice the volume of hydrochloric acid to reach the equivalence point - because of the 1 : 2 ratio in the equation. Suppose you start with 25 cm3 of sodium carbonate solution, and that both solutions have the same concentration of 1 mol dm-3. That means that you would expect the steep drop in the titration curve to come after you had added 50 cm3 of acid. The actual graph looks like this:

The graph is more complicated than you might think - and curious things happen during the titration. You expect carbonates to produce carbon dioxide when you add acids to them, but in the early stages of this titration, no carbon dioxide is given off at all. Then - as soon as you get past the half-way point in the titration - lots of carbon dioxide is suddenly released. The graph is showing two end points - one at a pH of 8.3 (little more than a point of inflexion), and a second at about pH 3.7. The reaction is obviously happening in two distinct parts. In the first part, complete at A in the diagram, the sodium carbonate is reacting with the acid to produce sodium hydrogencarbonate:

You can see that the reaction doesn't produce any carbon dioxide. In the second part, the sodium hydrogencarbonate produced goes on to react with more acid - giving off lots of CO2.

That reaction is finished at B on the graph. It is possible to pick up both of these end points by careful choice of indicator. That is explained on the separate page on indicators.

Adding sodium hydroxide solution to dilute ethanedioic acid Ethanedioic acid was previously known as oxalic acid. It is a diprotic acid, which means that it can give away 2 protons (hydrogen ions) to a base. Something which can only give away one (like HCl) is known as a monoprotic acid.

The reaction with sodium hydroxide takes place in two stages because one of the hydrogens is easier to remove than the other. The two successive reactions are:

If you run sodium hydroxide solution into ethanedioic acid solution, the pH curve shows the end points for both of these reactions.

The curve is for the reaction between sodium hydroxide and ethanedioic acid solutions of equal concentrations.