* Your assessment is very important for improving the workof artificial intelligence, which forms the content of this project
Download Topic 8: ACIDS and BASES
Chemical equilibrium wikipedia , lookup
Liquid–liquid extraction wikipedia , lookup
Peptide synthesis wikipedia , lookup
Electrochemistry wikipedia , lookup
Chemical reaction wikipedia , lookup
Hydrogen-bond catalysis wikipedia , lookup
Click chemistry wikipedia , lookup
Biological aspects of fluorine wikipedia , lookup
Physical organic chemistry wikipedia , lookup
Electrolysis of water wikipedia , lookup
Citric acid cycle wikipedia , lookup
Amino acid synthesis wikipedia , lookup
Nitric acid wikipedia , lookup
Acid throwing wikipedia , lookup
Sulfuric acid wikipedia , lookup
Metalloprotein wikipedia , lookup
Nitrocellulose wikipedia , lookup
Equilibrium chemistry wikipedia , lookup
Biosynthesis wikipedia , lookup
Fatty acid synthesis wikipedia , lookup
Butyric acid wikipedia , lookup
Biochemistry wikipedia , lookup
Lewis acid catalysis wikipedia , lookup
Nucleophilic acyl substitution wikipedia , lookup
Topic 8: ACIDS and BASES 8. 1. Theories of acids and bases 8.1.1 Define acids and bases according to the Brønsted–Lowry and Lewis theories. 8.1.2 Deduce whether or not a species could act as a Brønsted–Lowry and/or a Lewis acid or base. 8.1.3 Deduce the formula of the conjugate acid (or base) of any Brønsted–Lowry base (or acid). Bronsted-Lowry theory (applies in aqueous solutions) A Bronsted-Lowry acid is a proton donor A Bronsted-Lowry base is a proton acceptor Bronsted-Lowry acids must have a hydrogen/proton that can be released easily as an ion, H+ (a replaceable hydrogen). This happens when the hydrogen is part of a weak bond or polar bond. If the acid has 1 replaceable hydrogen it is called a monoprotic acid. The hydrogen ion released does not exist on its own in a solution as it is always accepted by another specie i.e. a base e.g. water to form an hydroxonium ion. HA (l) + (acid molecule) The above equation can also be written as: H2O (l) (base) HA (l) H3O+ (aq) H+ (aq) + + A- (aq) A- (aq) When acid molecules break up into ions in water we call it dissociation. Hydrogen ions are often called protons because they are H atoms, which have lost an electron. Examples of equations showing the dissociation/ionization of acid molecules. HCl + H2O HNO3 + H2O + H2SO4 + H2O + CH3COOH + H2O + H3O+ /H+(aq) + Cl-(aq) Other molecules are said to be acidic because they react with water to form an acid. e.g. sulfur trioxide is called an acidic oxide because of the reaction. SO3 + H2O H2SO4 Most non metal oxides are also acids because they act like acids i.e. they react with bases to produce a salt and water (neutralization reaction) Bronsted-Lowry bases are either substances which contain ions like OH-, O2- , HCO32- or CO32- as these ions can accept the proton/H+ (s) from the acid (because they contain atoms with 1 or more lone pairs); OH- (aq) + IB chemistry topic 8 H+ (aq) H2O (l) 1|Page 2H+ (aq) H2O (l) + CO2 (g) CO32- (s) + O2- (s) + H+ (aq) H2O (l) + CO2 (g) HCO3- (s) + H2O (l) 2H+ (aq) molecules that have lone pairs (e.g. ammonia and water); during an acid-base reaction, the base forms a co-ordinate sigma bond with the proton from the acid e.g. when ammonia and hydrogen chloride react to form ammonium chloride: NH3 + HCl NH4+ + ClIn this reaction the acid donates a proton to the base making the base a cation and the acid an anion; attraction between the two of them results in an ionic bond and the formation of NH4Cl. SO WHEN A BASE ACCEPTS A PROTON IT FORMS A DATIVE BOND WITH IT!!! According to the Bronsted-Lowry theory any reaction which involves the transfer of protons from one species to another is an acid-base reaction. Some substances can act as both and they are called amphoteric; what they behave like (either acid or base) depends on the other chemical in the reaction. For instance if the other specie is a stronger acid than our specie in question acts as a base, if the other specie is a stronger base than it will act as an acid. Conjugate acid-base pairs All Bronsted-Lowry acid-base reactions like…. the ionisation of acid molecules HCl (g) + H2O (l) we can also write this as H3O+ (aq) + HCl (aq) Cl- (aq) H+ (aq) + Cl- (aq) the ionisation of alkali molecules NH3 (g) + we can also write this as acid-base reactions H2O (l) NH4+ (aq) NH4OH (aq) HCl (aq) + NH3 (aq) + OH- (aq) NH4 + (aq) + NH4+ (aq) OH- (aq) + Cl- (aq) are reversible processes (which means an equilibrium exists). This is the case even when it involves strong acid or strong bases like hydrochloric acid and sodium hydroxide. In the case of strong acids and bases the reverse reaction is ignored. Within a reversible system, just like the forward reaction, the reverse reaction can also be considered a Bronsted-Lowry acid-base reaction as it also involves a transfer of protons but in the opposite direction. Therefore, in the equilibrium process there are two acids and two bases; an acid and a base on each side of the equation. These acids and bases in such an equilibrium process are related: the acid on the reactant side is related to the base (this base is called the conjugate base of the acid) on the product side in the following way: IB chemistry topic 8 2|Page An acid + its conjugate base (= acid - H+ ) = conjugate acid-base pair A base + its conjugate acid (= base + H+) = conjugate acid-base pair An acid dissociates into a proton and its conjugate base which is its anion. The base on the reactant side is related in the same way to the acid on the product side; this acid is called the conjugate acid of the base. Example: HCl (aq) + acid NH3 (aq) NH4+ (aq) conjugate acid of the base NH3 Conjugate acid-base pairs from the above example: Cl- (aq) + base conjugate base of the acid HCl HCl and ClNH3 and NH4+ . A conjugate acid-base pair is a pair of species which differ by the presence of one H+ only; the acid member of the pair always has an extra H+. When an acid donates a proton it becomes a conjugate base. When a base accepts a proton it becomes a conjugate acid. A conjugate base behaves like a base while a conjugate acid behaves like an acid in the reverse reaction. As stated earlier some substances can behave like both acids and bases depending on the reaction they are involved in; they are referred to as amphotheric substances e.g. water, HSO 4- and HCO3-. Exercises 1. In which one of the following reactions is the species in bold type behaving as a base? A. 2NO + O2 2NO2 C. NH4+ + H2O NH3 + H3O+ B. CO32- + H+ HCO3D. Cu2+ + 2OH- Cu(OH)2 2. Which one of the following is the conjugate base of the hydrogen sulfite ion, HSO3- ? B. H2SO3- A. H2SO3 C. SO32 - D. SO3- 3. What is the conjugate acid of: CH3COO-, HSO4-, NH3 , OH- , F- ? 4. What is the conjugate base of: HCl, H3O+, HSO4-, NH3 ? 5. Which one of the following species, many of which are unstable, would you expect to be capable of acting as a base? B. CH3 A. CH4 C. CH3+ D. CH3- 6. Which one can function as both an acid and a base according to the Bronsted-Lowry definition? A. HS- B. S2- C. NH4+ D. Al3+ 7. According to the equations below, what is the conjugate acid of HPO42H3PO4 + H2O H2PO4- + H3O+ H2PO4- + H2O HPO42- + H3O+ HPO42- + H2O PO43- + H3O+ A. H3PO4 IB chemistry topic 8 B. H2PO4- C. HPO42- D. PO433|Page 8. All the following are conjugate acid-base pairs EXCEPT A. OH- , O2- B. H2CO3, HCO3- C. HNO2, NO2+ D. HSO4- , SO42- 9. Which of the following is NOT an acid-base pair? A. HNO3 / NO3- B. NH3 / NH2- C. H2SO4 / HSO4- D. H3O+ /OH- 10. In each of the following equations identify the conjugate acid-base pairs (a) HNO2 + H2O H3O+ + NO2(b) SO42- + H3O+ H2O + HSO4(c) H- + H2O H2 + OH(d) H3O+ + OH- 2H2O (e) HSO3- + H2O H2SO3 + OH(f) H2 C2O4 + H2O HC2O4- + H3O+ Structures of conjugate acid-base pairs Members of a conjugate acid-base pair always differ by a single proton. To recognise a conjugate acid or base more quickly, their written structures should always make clear the approximate location of the proton that is transferred. This applies especially to organic conjugate acid-base pairs. Examples: ethanoic acid: CH3COOH and CH3COO ethanol: C2H5OH and C2H5O phenol: C6H6OH and C6H6ORelationship between acid and conjugate base and base and conjugate acid If an acid is strong its conjugate base is weak; the stronger the acid the weaker its conjugate base; the more the equilibrium lies towards the products. The weaker the acid, the stronger the conjugate base which means the reverse reaction is favoured and the equilibrium lies towards the reactants. The stronger the base the weaker its conjugate acid; the weaker the base, the stronger its conjugate acid. Lewis Theory: a much broader theory Using this theory, a lot more reactions can now be described as ‘acid-base’ reactions. All Bronsted/Lowry acid-base reactions are also Lewis acid-base reactions! Lewis acids: electron pair acceptors A Lewis acid is a reactant which during a reaction accepts an electron pair from another substance; Common Lewis acids are: boron and aluminium compounds in which neither element obeys the octet rule but has empty orbitals in which it can place an electron pair; IB chemistry topic 8 4|Page positive ions like the H+ which never exists on its own and either reacts with water to form H3O+ or with a another base to form water. Other examples include CH3+ and Br+ (Hal+) which you will study in organic chemistry. molecules containing positive centres (mostly organic molecules) as a result of polar bonds within the molecule e.g. C in CO2 or in halogenalkanes and S in SO2. All Bronsted-Lowry acids are Lewis acids but not vice versa; Although strictly speaking dissociation needs to occur first before a Bronsted acid like HCl can act as a Lewis acid; it is only the H+ which can accept the electron pair and therefore behave like a Lewis acid. BF3 can act as a Lewis acid but not a Bronsted-Lowry. Lewis acids are electron deficient and are looking to gain electrons. Lewis bases: electron pair donors A Lewis base is a reactant which during a reaction donates an electron pair; common Lewis bases are: nitrogen compounds (e.g. ammonia), water and the hydroxide ion as they have a great tendency, because of their lone pairs, to donate these electron pairs to Lewis acids. Also all Bronsted-Lowry bases are also Lewis bases because the Bronsted bases can only accept protons by donating an electron pair to them. HIGHER LEVEL: Ligands (neutral molecules or anions with lone pairs) are Lewis bases that attach themselves to metal ions which act like Lewis acids e.g. hydration of transition metal ions. CuSO4.5H2O is in reality [Cu(H2O)4]2+ SO42- .H2O; Lewis acid-base reactions A Lewis acid-base reaction involves the formation of a dative bond; a covalent bond in which both electrons (lone pairs) are donated by the base to the acid; Examples: H+ + OH H2O NH3 + BF3 NH3BF3 Exercises 1. For each of the following species, state whether it is most likely to behave as a Lewis acid or Lewis base. Explain your answer. a) PH3 b) BCl3 c) H2S e) Cu2+ d) SF4 2. All of the following species can function as Lewis bases EXCEPT A. OH- B. N3- D. NH4+ C. NH3 3. Classify each of the following species as a Lewis acid or base. (a) CO2 (b) H2O (c) I- (d) SO2 (e) NH3 (f) OH- (g) BCl3 (h) NH4+ (i) H+ (j) N3- 4. Identify the Lewis acid and base in the following reactions. (a) Hg2+ + 4CN- Hg(CN)42- IB chemistry topic 8 (e) CN- + H2O HCN + OH- 5|Page (b) SnCl4 + 2Cl- SnCl62(c) Co3+ + 6NH3 Co(NH3)63+ (d) CH3COO- + 2HF CH3COOH (f) SO2 + H2O H2SO4 (g) HNO2 + OH- NO2- + H2O + HF2- 8. 2. Properties of acids and bases 8.2.1 Outline the characteristic properties of acids and bases in aqueous solution. Acids For an acid to show its acidic properties it must react with a base e.g. water. During that reaction the acid molecule which has a covalently bonded (weak covalent bond or polar covalent bond) hydrogen atom in it, splits into two ions: H+ and a negative non-metal ion like eg Cl- or NO3- according to the following general equation: HA + H2O H3O+ + A- HA or H+ + A- This reaction is a disscociation or ionisation reaction. Within the scope of the syllabus we will limit our study to monoprotic acids that are acids with 1 hydrogen that they can donate. Examples of diprotic acids are sulfuric acid and carbonic acid. It is the H3O+/H+ ions which give acids their properties: effect of acids on indicators: * blue litmus turns red; * phenolphthalein remains colourless; * methyl orange: red-orange for low pH; * universal indicator: deep red to red orange; pH 7; react with metal oxides and hydroxides to form salts and water (=neutralisation); examples: acid + metal oxide: acid + metal oxide: acid + metal hydroxide: CuO(s) + 2HCl (aq) CuCl2 (aq) + H2O (l) MgO(s) + 2CH3COOH (aq) Mg(CH3COO)2 (aq) + 2NaOH(s) + H2SO4 (aq) Na2SO4 (aq) + H2O (l) 2H2O (l) the above neutralisation reaction can be represented by the net ionic equation: 2OH- (aq) + 2H+ (aq) Mg (s) Mg(NO3)2 (aq) + H2 (g) react with metal carbonates and metal hydrogen carbonates to produce a salt, carbon dioxide and water; examples: CaCO3 (s) + 2HCl (aq) CaCl2 (aq) + H2O (l) + CO2 (g) 2NaHCO3 (s) + H2SO4 (aq) (net ionic equation for the first one = as the sodium and sulfate ions are spectator ions. react with the more reactive metals to produce a salt and hydrogen; 2HNO3 (aq) + 2H2O (l) Na2SO4 (aq) + 2H2O (l) + CO2 (g) CO32- (s) + 2H+ (aq) CO2 (g) + 2H2O (l)) react with ammonia to form ammonium salts, e.g.: NH3 (s) + HCl (aq) IB chemistry topic 8 NH4Cl (s) or 2NH3 (s) + H2SO4 (aq) (NH4)2 SO4 (aq) 6|Page electrolytes as they ionise as a result of the dissolution caused by a polar solvent which in most cases will be water; the greater the ion concentration the greater their conductivity. Bases metal oxides, metal hydroxides, metal carbonates, metal hydrogen carbonates, amines; ALKALIS = bases which are soluble in water; they form the OH- ion when they dissolve in water (they either have the hydroxide ion as part of the compound and just dissociate – e.g. NaOH – or they produce it when they react with water e.g. Li2O, NH3 or CO32-) . Na+ (aq) + OH- (aq) NaOH (aq) Li2O (s) + H2O (l) NH4 + (aq) + OH- (aq) NH3 (g) + H2O (l) CO32- + H2O (l) effects on indicators: red litmus turns blue; phenolphthalein: pink; 2Li+ (aq) + 2OH- (aq) HCO3- (aq) + OH- (aq) methyl orange: yellow; alkalis: universal indicator: green to deep violet; pH 7; react with and neutralise acids to form salts and water; alkalis are electrolytes as they dissociate as a result of the dissolution caused by a polar solvent like water. Exercise: For each of the following chemical reactions write a balanced symbol equation and an ionic equation. 1. 2. 3. 4. 5. 6. 7. magnesium and sulfuric acid sodium hydrogen carbonate and hydrochloric acid sulfuric acid and sodium chloride calcium hydoxide and hydrochloric acid copper and nitric acid lead carbonate and nitric acid iron with dilute sulphuric acid 8. 3. Strong and weak Bronsted acids and bases 8.3.1 Distinguish between strong and weak acids and bases in terms of the extent of dissociation, reaction with water and electrical conductivity. 8.3.2. State whether a given acid or base is strong or weak. 8.3.3 Distinguish between strong and weak acids and bases, and determine the relative strengths of acids and bases, using experimental data. Strength of an acid Strong acids ionize or dissociate (almost) completely (=equilibrium process which goes to virtual completion) when they react with water. In weak acids not all of the acid molecules ionize, in fact most acid molecules remain molecules; the weaker the acid the fewer the number of acid molecules that ionize. The weaker the acid the more the equilibrium of ionisation lies towards the reactants; the lower the concentration of hydrogen ions in solution. When a strong acid is added to water the ionisation is not really represented by an equilibrium: e.g.: IB chemistry topic 8 HNO3 H+ + NO3- 7|Page !!!!! We assume that all the H+ in the solution come from the acid and not the water which also dissociates but only very slightly. The strength of an acid is controlled by: how easy it is to break the bond joining the hydrogen to the rest of the molecule; if the bond is weak, the acid is strong. For example, although a covalently bonded molecule and very polar, HF is a weak acid because the covalent bond is strong as a result of the small size of both the H and F atoms. HI is a stronger acid than HBr and HCl as its hydrogen-halogen bond is the weakest. how stable the anion is that is formed when the acid molecule donates its proton eg NO3-. If the anion is stable the acid is strong as the anion does not readily react with a proton to form the acid molecule again - it is a weak conjugate base; the nature of the parent molecule eg ethanol is a weaker acid than ethanoic acid (more of this in organics) Therefore, a strong acid ionizes readily or gives up very readily its hydrogen (either to water when it dissolves or to a base). Because of their different extent of ionization, weak acids have a lower concentration of H+ in their aqueous solutions, and therefore a higher pH, than strong acids of the same concentration. Strong acids: HCl, HNO3, H2SO4 , HBr, HI, H3PO4. Weak acids: CH3COOH (=ethanoic acid), H2CO3 (carbonic acid), HCOOH, citric acid, all carboxylic/organic acids. Exercise: For each of the above acids write an equation showing their dissociation in water. Strength of bases Strong bases dissociate (almost) completely (i.e. their ions separate) when they dissolve in water while weak bases do not dissociate completely. A strong base accepts a proton very readily. NaOH (s) e.g.: OH- (aq) + Na+ (aq) Weak bases have a lower concentration of OH- and therefore a lower pH than strong bases of the same concentration. Most of their particles/molecules have not accepted protons. Strong bases: all group 1 hydroxides and Ba(OH) 2 (all hydroxide ions in the solution are considered to be coming from the base and not water); Ba(OH) 2 (s) 2OH- (aq) + Ba2+ (aq) Weak bases: NH3 and amines. NH3 (g) + H2O (l) NH4+ (aq) + OH- (aq) Difference between strength and concentration. the concentration of an acid or base refers to the number of acid or alkali molecules that are present in 1L of the solution; when an acid is diluted more water is added so that there are fewer acid molecules per 1L of solution; strength refers to the degree of ionisation or dissociation that the acid or base undergoes when dissolved in water. IB chemistry topic 8 8|Page Example: a hydrochloric acid solution and an ethanoic acid solution of the same concentration (eg 0.1M) will have different hydrogen concentrations - and therefore different pH’s - because of their different strengths: the pH of the hydrochloric acid will be lower than the pH of the ethanoic acid solution. Some practical procedures to distinguish between strong and weak acids and bases. Simple procedures to determine relative acidities or basicity of substances: It is important that in all these tests, the acids or alkalis which are being compared are of the same concentration. All measurements taken in the procedures below are dependent on the amount of H+ ions present in the solution. For example, when comparing two 1 mol dm-3 solutions of H2SO4 and HCl, there are twice as many H+ ions in HsSO4 as in HCl so therefore the H2SO4 conducts better. comparison of pH of substances; compare pH with known strong acids; if the pH is higher the acid is weak. comparison of conductivity of same concentrations; measure the conductivity of solutions of the same concentration. The acidic or basic solutions with the largest concentration of ions will have the highest conductivity as those with the highest conductivity have dissociated more. compare rate of reactions with reactive metals; e.g. measure gas production; strong acids react at higher rates than weaker acids. 8. 4. The pH scale 8.4.1 Distinguish between aqueous solutions that are acidic, neutral or alkaline using the pH scale. 8.4.2 Identify which of two or more aqueous solutions is more acidic or alkaline using pH values. 8.4.3 State that each change of one pH unit represents a 10-fold change in the hydrogen ion concentration [H+(aq)]. The pH scale changes is used to comparewhen neutrality, or basicity of by different solutions. 8.4.4Deduce in [H+(aq)] the pH acidity of a solution changes more than one pH unit. The pH is measured using a pH meter or pH paper; the latter consists of a mixture of indicators which respond by showing different colours at different concentrations of H+ ; each indicator shows a particular colour at a certain concentration of H+. The lower the pH the greater the acidity; the greater the pH, the greater basicity and the lower the concentration of H+. The pH can be derived from the concentration of the H+. As the [H+] is often a very small number the negative logarithm is taken; a negative logarithm to end up with a positive number. So: pH = - log [H+] pOH = - log [OH-] Examples: an acidic solution with a hydrogen ion concentration of 1 x 10-3 has a pH of 3. an alkali solution with a [H+] of 1 x 10-9 has a pH of 9. So one pH unit represents a tenfold change in acidity or basicity !!!! Example: a solution with a pH of 2 has: a tenfold increase in [H+] than a solution with a pH of 3 100 times more [H+] than a solution with a pH of 4 but a tenfold decrease in [H+] than s solution with pH 1. IB chemistry topic 8 9|Page Use the following simulations: http://phet.colorado.edu/en/simulation/acid-base-solutions to reinforce ideas about strong and weak acids and alkalis and pH. http://www.chem.iastate.edu/group/Greenbowe/sections/projectfolder/flashfiles/acidbasepH/ph_ meter.html http://phet.colorado.edu/en/simulation/ph-scale http://www.afn.org/~afn02809/#ACID-BASE%20CHEMISTRY%20RESOURCES IB chemistry topic 8 10 | P a g e