Aims to increases students understanding of: History, nature and practice of chemistry Applications and uses of chemistry Implications of chemistry for society and the environment
1. Definitions: Lavoisier / Davy; Arrhenius; Brønsted- Lowry; conjugate pairs; salts acidic/basic/neutral; amphoteric substances; neutralisation; titrations; standard solutions; ph; buffers 2. Indicators: natural, common lab-based; uses of outside the lab; identification of acidic, basic, neutral; Arrhenius definition
3. Acids in the air: non-metal oxides (e.g. CO 2, NO, NO 2, SO 3 ) as acids; equilibria; Le Chatelier s Principle; natural and industrial sources, chemical equations; assessment of evidence; environmental consequences 4. Metal oxides (e.g. CuO, MgO, CaO, Fe 2 O 3 ) as bases; Periodic Table; solubility; 5. Acids in Biological Systems: Brønsted-Lowry definitions; simple organic acids; strong/weakdilute/concentrated distinctions; degree of ionisation weak/strong/equilibria
Antoine Lavoisier and Humphry Davy Lavoisier thought acids contained oxygen as lots of oxides are acidic (but many are not!) Davy thought acids contained hydrogen (as replaceable hydrogen Hydrogen that could be replaced with a metal) 2004 and 2010 exams asks for the Davy definition of an acid
Arrhenius Definition of Acids and Bases An acid is a substance that provides H + ions* in aqueous solution. e.g. HCl(g) H + (aq) + Cl (aq) e.g. H 2 SO 4 (l) 2H + (aq) + SO 4 2 (aq) *Note do not have free H + ions as it is attached to water to form the hydronium ion H 3 O + H H O + H + + O H H formation of hydronium ion H
Arrhenius definition of base A base is a substance that provides OH (hydroxide) ions in aqueous solution. e.g. NaOH(s) Na + (aq) + OH (aq)
The hydronium ion becomes surrounded by water. Therefore generally refer to ionic species as aquated species: e.g. HCl(g)+H 2 O(l) H 3 O + (aq) + Cl (aq) For convenience commonly write: HCl(g) +H 2 O(l) H + (aq)+ Cl (aq)
This applies to all ions in water NaCl in water looks like:
Problems with Arrhenius s Definition The Arrhenius definition explains the properties of many acids and bases in the presence of water, but not all. For example: Many species act as acids/bases in solvents other than water And even in the absence of ANY SOLVENT Their reactivity is dependent on the solvent and/or other species present Q8 in the 2008 paper asks about Arrhenius definition of acids
Brønsted Lowry Definition of Acids and Bases An acid is a proton (H + ) donor molecule or ion A base is a proton acceptor molecule or ion H acid base proton transfer H Acid base reactions occur simultaneously
H acid base proton transfer H e.g. HCl(g) + NH 3 (g) NH 4 Cl(s) NH 4 Cl => made up of NH 4+ and Cl Hydrogen chloride is a Brønsted Lowry acid as it donates a proton Ammonia is a Brønsted Lowry base because it accepts a proton
H 2 O(l) + NH 3 (g) NH 4+ (aq) + OH - (aq) H 2 O(l) + NH 4+ (aq) H 3 O + (aq) + NH 3 (aq) charge is always conserved the sum of charges on each side is the same 0 + 0 = 1 + (-1) 0 + 1 = 1 + 0 mass is always conserved the number of atoms of each species on each side is the same 2H + 3H, O, N = N, O, 4H + H 2H + 4H, O, N = N, O, 3H + 3H There was a question in 2002 HSC on Brønsted-Lowry acids
Summary Lavoisier - noticed that many oxides were acidic (eg nitrous oxide, sulfur trioxide) Davy - noticed many hydrides were acidic (hydrogen sulfide, hydrogen chloride) Arrhenius - acids dissociate to H + and bases to OH Brønsted/Lowry - acids are are, bases 2005 HSC had the question Analyse how knowledge of the composition and properties of acids has led to changes in the definition of acids.
Conjugate Pairs An acid gives up a proton to form its conjugate base A base accepts a proton to form its conjugate acid This is what forms a BUFFER (Q9 2005 and Q6 H proton transfer 2008 HSC) H acid base conjugate base conjugate acid
Base ----------------------- Conjugate acid e.g. NH 3 (l) + H 2 O(l) NH 4+ (aq) + OH (aq) Acid ----------------------Conjugate base Conjugate base = acid H + Conjugate acid = base + H + Remember charges must balance on both sides of the equation! Acid ------------------------- Conjugate base e.g. NH 3 (l) + OCH 3 (aq) NH 2 (aq)+ CH 3 OH(l) Base ------------------------- Conjugate acid
The conjugate base of a strong acid is an extremely weak base (e.g.hcl Cl ) The conjugate base of a weak acid is a weak base (e.g. NH 4+ NH 3 ) The conjugate bases of extremely weak acids (e.g. water, methanol and ethanol) are very strong bases Q7, 2009 asked about the conjugate base of HSO 4
A Closer Look: Electron Movement H 2 O HCl H 3 O Cl As Lewis structures, with all valence electrons shown: How does the proton change positions? Curved arrows are used to show the movement of electron pairs O-H bond is being made H H O H H Cl H O H Cl H-Cl bond is being broken
A Closer Look: Electron Movement As Lewis structures, with all valence electrons shown: How does the proton change positions? N-H bond is being made H-Cl bond is being broken Which is generally written as:
Strong and Weak Acids and Bases Strong acids and bases are completely ionised. Weak acids and bases are partially ionised. H 2 SO 4, HCl, HBr, HI and HNO 3 are strong acids. Most other acids (e.g. ethanoic acid, citric acid, NH 4 + and carbonic acid) are weak acids. Metallic oxides and hydroxides are generally strong bases (remember Lavoisier?). Ammonia, ethanoate (acetate) and carbonate ions are examples of weak bases.
Ionisation of any acid (HA) in water: HA(aq) + H 2 O(l) H 3 O + (aq)+ A (aq) The degree of ionisation = [H 3 O + ] 100% [HA] Ionisation of base (B) in water: B(aq) + H 2 O(l) HB + (aq)+ OH(aq) The degree of ionisation = [ OH] 100% [B] [H 3 O + ] = concentration of H 3 O + produced, get from ph [HA] = concentration as made up, get from moles of acid in volume of solution
How much does water ionize? Autoionisation of water: H 2 O(l) + H 2 O(l) H 3 O + (aq) + OH (aq) K W [H 3 O ] [OH ] K w = 10 14 [H 3 O + ].[OH ] = 10 14 At neutrality [H 3 O + ] = [OH ] Therefore [H 3 O + ] = 10 7 M ph = log 10 [H 3 O + ] = 7
The ph scale and ph calculations ph = log 10 [H 3 O + ] A neutral solution has a ph of 7 (exactly) An acidic solution has a ph <7 ([H 3 O + ] > [ OH]) A basic solution has a ph >7 ([ OH] > [H 3 O + ]) At 25 C: [HO ] = 1.00 10 14 OR poh = 14 ph [H 3 O + ]
ph Calculations ph of strong acid/base calculated from molarity The ph of a weak acid/base requires the degree of ionisation. Example: Calculate the ph of a 0.1 molar solution of ascorbic acid. Note the solution is 2.8% ionised. i) Degree of ionisation = [H 3 O + ] 100% = 2.8% [ascorbic acid] ii) 2.8 = [H 3 O + ] 100% 0.1 iii) [H 3 O + ] = 2.8 10 3 molar iv) ph = log 10 [2.8 10 3 ] = 2.6 Exams generally have a question on calculation of ph. e.g. Q8 2005; Q10 2007; Q21 2008; Q21 2010 HSC exam
Be practical. A solution of a strong acid (100% ionised) has a lower ph value (i.e. higher [H 3 O + ]) than the ph of a similar concentrated solution of a weak acid (not fully ionised). See hydrochloric acid verses ascorbic acid in 2001 exam (Q20), 2002 (Q22) and hydrochloric verses ethanoic (acetic) acid in 2010 (Q21).
Generally the reaction of an acid and base produce a salt and water. e.g. HCl(aq) + NaOH(aq) NaCl(aq) + H 2 O(l) The reaction of an acid and base to give salt and water is often referred to as a neutralisation reaction. The reaction is exothermic, i.e. releases heat.
Buffers: Le Chatelier s Principle Buffers are mixtures of a weak acid and weak base - usually conjugate pairs Buffer solutions buffer against changes in ph Both components are weak and readily accept protons added to the solution if they are a base or have protons to donate to added bases if they are an acid - Le Chatelier s Principle Acetic acid & sodium acetate CH 3 COOH + B CH 3 COO - + BH + CH 3 COO - + HA CH 3 COOH + A -
Le Chatelier s Principle in action http://www.mhhe.com/physsci/chemistry/essentialchemistry/flash/buffer12.swf
Natural / Biological Buffer Systems Important biological buffer systems dihydrogen phosphate system internal fluid of all cells H 2 PO 4 (aq) H + (aq) + HPO 2 4 (aq) carbonic acid system sea water and blood CO 2 (g) + H 2 O(l) H 2 CO 3 (aq) H + (aq) + HCO 3 (aq) ph of ocean has dropped from 8.1 to 7.9 in the last 20 years Both bicarbonate and hydrogenphosphate are amphiprotic (a substance that can donate or accept a proton, H + ) This was asked in 2004 and 2006 HSC
ph of Salt Solutions A salt solution may be acidic, basic or neutral Depends on strength of parent acid and base Strong Acid + Strong Base = Neutral Salt Strong Acid + Weak Base = Acidic Salt Weak Acid + Strong Base = Basic Salt HCl + NaOH NaCl + H 2 O Neutral HCl+ NH 3 NH 4 Cl Acidic NaOH + CH 3 COOH CH 3 COONa + H 2 O Basic
Volumetric Analysis of Acids and Bases Used to determine concentration of unknown acid/base Equivalence Point: Where amounts of acid & base sufficient to cause complete consumption of both. Measure equivalence point with indicator, ph meter or by electrical conductivity.
For example: When 20 ml of 0.1 M hydrochloric acid is titrated with 0.1 M sodium hydroxide (from a burette) the following conductance change is observed. Reaction: H + + Cl + Na + + HO H 2 O + Na + + Cl As NaOH is added, the H + gets used up and is effectively replaced by Na + ions, which are much larger than H + ions and have a smaller conductivity. Conductance falls as the equivalence point is reached. After the equivalence point, we are essentially adding NaOH to the NaCl solution. This adds extra ions and therefore leads to an increase in conductance. Conductivity 0 10 20 30 ml NaOH (aq) added *Understanding of a conductivity graph was assessed in the 2001 exam.
Example of a Volumetric Acid Base Analyses From 2008 HSC exam, Q28: A standard solution was prepared by dissolving 1.314 g of sodium carbonate in water. The solution was made up to a final volume of 250.0 ml (a) Calculate the concentration of the sodium carbonate solution Sodium carbonate is Na 2 CO 3, MW = 2 22.99 + 12.01 + 3 16.00 = 106.0 g mol -1 (4 sig. fig.) Na 2 CO 3 = 1.314 g / 106.0 g mol -1 = 0.01240 mol c = 0.01240 mol / 0.2500 L = 4.959 10-2 mol L -1 molarity is a measure of concentration molarity = no. moles volume (in L) M = m V no. moles = mass in g (that we have) MW (molecular weight)
Example of a Volumetric Acid Base Analyses This solution was used to determine the concentration of a solution of hydrochloric acid. Four 25.00 ml samples of the acid were titrated with the sodium carbonate solution. The average titration volume required to reach the end point was 23.45 ml. (b) Write a balanced equation for the titration reaction Na 2 CO 3 (aq) + 2 HCl(aq) 2 NaCl(aq) + CO 2 (g) + H 2 O(l) acid + carbonate salt + carbon dioxide + water (c) Calculate the concentration of the hydrochloric acid solution Na 2 CO 3 = 4.959 10-2 mol L -1 0.02345 L = 1.163 10-3 mol molarity = no. moles volume (in L) no. moles = molarity x volume (in L)
Example of a Volumetric Acid Base Analyses This solution was used to determine the concentration of a solution of hydrochloric acid. Four 25.00 ml samples of the acid were titrated with the sodium carbonate solution. The average titration volume required to reach the end point was 23.45 ml. Na 2 CO 3 (aq) + 2 HCl(aq) 2 NaCl(aq) + CO 2 (g) + H 2 O(l) (c) Calculate the concentration of the hydrochloric acid solution from the reaction stoichiometry, we need twice as many moles of HCl compared with Na 2 CO 3 moles of HCl =2 1.163 10-3 mol = 2.326 10-3 mol [HCl] = 2.326 10-3 mol / 0.02500 L = 9.303 10-2 mol L -1 molarity = no. moles volume (in L) M = m V
(Another) Example of a Volumetric Acid Base Analyses From 2001 HSC exam: A household cleaning agent contains a weak base of general formula NaX. 1.00 g of this compound was dissolved in 100.0 ml of water. A 20.0 ml sample of the solution was titrated with 0.1000 mol L 1 hydrochloric acid, and required 24.4 ml of the acid for neutralisation. What is the molar mass of this base? Step 1: write a balanced equation: NaX(aq) + HCl(aq) NaCl(aq) + HX(aq) for every mole of NaX ONE mole of HCl is needed.
Step 2: calculate moles of HCl used for the titration: = 0.1000 mol L 1 24.4 10 3 L = 2.44 10 3 moles (3 sig. figures) = number of moles NaX in the 20.0 ml sample titrated. molarity = no. moles volume (in L) no. moles = molarity x volume (in L) Step 3: calculate moles of NaX in original 100 ml and therefore 1.00 g: = 100/20 (taking into account only analysed 20 ml) 2.44 10 3 = 1.22 10 2 moles we only titrated 20 ml of the original 100 ml, i.e., we only used 1/5 of the initial amount, therefore the number of moles in 1 g is 5 x our answer from above Step 4: calculate molar mass of NaX: Remember moles = 1.22 10 2 moles = mass (1.00 g)/molar mass molar mass = 1.00 g/1.22 10 2 moles = 82.0 g/mole (3 sig. figures) no. moles = mass in g mass in g MW x no.moles = mass in g MW = MW no. moles Q23 in 2003, Q24 in 2005, Q21 in the 2007 and Q28 in 2010 HSC exams. For full marks, show all steps; including providing a balanced equation, volumetric analyses questions ALWAYS are assessed!
Indicators for Acid Base Titrations ph ~7 Bromothymol blue colour change ph 6.2 7.6 ph > 7 Phenolphthalein colour change ph 8.3 10 Br H 3 C HO O HO H 3 C O S O CH 3 CH 3 CH 3 OH Br CH 3 OH O ph < 7 Methyl orange O colour change ph 4.4 3.1 H 3 C N H 3 C N N S O O OH There was a question in 2007, 2009 & 2010 HSC on selection of indicators and in 2003 was to identify what indicator was used.
Natural indicators Red (Purple) Cabbage; Red cabbage contains a mixture of pigments used to indicate a wide ph range. Red onion also changes from pale red in an acidic solution to green in a basic solution Tea; Tea is an acid-base indicator. Its colour changes from brown in basic solution to yellow-orange in acid. This is the reason why lemon juice is added to a cup of tea not for the taste, but to reduce the intensity of the colour (called brightening ) Also beetroots, cherries, many berries and grapes contain ph sensitive pigments
Non-metal and metal oxides CaO(s) + H 2 O(l) Na 2 O(s) + H 2 O(l) Ca(OH) 2 (aq) 2 NaOH(aq) but SO 3 (g) + H 2 O(l) H 2 SO 4 (aq) 2 NO 2 (g) + H 2 O(l) HNO 2 (aq) + HNO 3 (aq) Questions about the acidity/basicity of oxides was asked in 2002 and 2009 HSC