Module 6: acid/base reactions

Acid + metal → salt + hydrogen

(not Bronsted-Lowry acid-base reaction, instead a redox reaction). Exothermic, strength depends on concentration and strength of acid and position of metal on activity series, use pop test.

Effect of dilution on the pH of strong acids and bases

(often diluted for desired concentration) - pH increases when a solution of acid is diluted (as lower concentration of H+ ions in relation to water). **However, cannot increase past 7, as dilution does not change the properties of an acid to accept protons, it still donates and will remain strong but be dilute. pH in this way not necessarily an indication of strength, need to know concentration as well. Similarly, dilution decreases the pH of a base (as lower conc of OH-), though again can't go below 7.

Enthalpy of neutralisation

(ΔHneut) - heat released per mol of water produced during a neutralisation reaction, kJ/mol. Neutralisation is exothermic so ΔHneut <0. When strong acid react with strong bases, the hydrogen and hydroxide ions that react to form water are already within solution. Using calorimetry, q=mcΔT (c=4.18 as per usual), then ΔHneut = -q/n(H2O).

Bronsted-Lowry

*Acids are proton donors, bases are proton acceptors. Acid base reactions involve exchange of protons from an acid to a base. More general, substances are acidic when they donate a proton H+ to a base, substances are basic if they accept a proton from an acid. Advantages: explains solvents and solutions in different contexts, not just in aqueous solutions. Can be in gases (e.g. HCl (g) + NH3(g) → NH4Cl(s)). Explains ammonia as a base, which accepts a proton to become ammonium. Limitations: still requires to have an H+ ion to donate. E.g. HCl(g) + H2O(l) → H3O+(aq) + Cl- (aq) - HCl donated proton to water to form hydronium, so HCl is an acid. Water has accepted a proton from the HCl so water is a base.

Monoprotic base

Accepts one proton e.g. NH3.

Davy (Davy displacement)

Acids are acidic due to presence of replaceable hydrogen, found after producing hydrogen gas when reacting acids with metals. Also found acid + base = salt + water Limitations: not all acids dissociate to hydrogen (ammonia, same issues as Arrhenius).

Lavoisier

Acids are acidic due to replaceable oxygen, applies to H2SO4, HNO3 and H3PO4. Limitations: not all acids have oxygen, e.g. non metal oxides HCl.

Arrhenius

Acids dissociate and ionise in water to produce H+ (or H3O+) ions, bases dissociate to produce OH- ions. Advantages: simplified definition, technically correct (although a generalisation), more easily differentiates between acids and bases (H+ ions vs. OH- ions), correct in most contexts (aqueous). Limitations: only worked in aqueous environments (so no gases or solvents other than water), couldn't explain ammonia (which has no hydrogen or hydroxide), couldn't explain how not all substances are acidic or basic in different solvents, couldn't explain the basicity of oxides and carbonates.

Molarity (mol/L)

C=n/V C1V1=C2V2 , you know the drill, dilution doesn't change the amount of an acid or base rather changes the volume and concentration. Can prepare appropriate solutions of a base by: accurately weighing out a mass of the base, transferring base to a volumetric flask, ensuring complete transfer of the base by washing with water, dissolve the base in water, adding water to make the solution up to the calibration mark and shake thoroughly.

Weak acids

CH3COOH, H2CO3, H3PO4, HF, C6H8O7 (citric acid) Weak acids prefer to stay as molecules, only partially dissociating (thus equilibrium reaction, equilibrium lies to left). <100% ionisation.

Polyprotic acids

Can donate more than one proton (duh), they donate in two stages when reacting with a base. Are reversible reactions.

Monoprotic acids

Can donate only one proton e.g. HCl, HF, HNO3, CH3COOH (only the hydrogen that is part of OH is donated so ethanoic acid is monoprotic).

Bases

Caustic, bitter, slippery, react with fats to produce water soluble soaps, neutralise acids, pH>7, solutions conduct electric current.

Naming acids

Change 'ide' to 'ic' for non-oxyacids e.g. hydrogen chloride to hydrochloric acid. Change 'ate' to 'ic' for oxyacids e.g. anion sulfate SO42- to sulfuric acid H2SO4. Change 'ite' to 'ous' for oxyacids e.g. anion sulfite SO32- to sulfurous acid H2SO3.

Phenolphthalein

Colourless Colourless Pink - crimson Good to differentiate strongly and weakly alkaline. Does not lose protons as easily. pH 8.3 - 10

Acids

Corrosive, sour, neutralise bases, pH<7, solutions conduct electric current, turn blue litmus red, react with and dissolve active metals, react with carbonates

Triprotic acids

Donate three protons e.g. H3PO4, H3BO3 and C6H8O7 (ion C6H5O73-, citric acid), they donate in three stages. Phosphoric acid a weak acid in water, the extent of dissociation decreases progressively from stage 1 to stage 3.

Neutralising excess stomach acidity

Effervescent NaHCO3 antacid tablets treats symptoms of acid reflux and indigestion 3NaHCO3 + C3H5O(COOH)3 → Na3C3H5O(COO)3 + 3H2O + 3CO2

Amphoteric

Gains or loses, acting as an acid or base e.g. could be zinc losing electrons, not necessarily protons.

Strong acids

HCl, H2SO4, HNO3, HClO4 (perchloric acid), Strong acids readily donate protons, thus solutions of strong acids contain ions with virtually no molecules, ionising completely (100%). Tendency to lose/ dissociate protons.

Water

Has amphiprotic properties, behaving as a very weak acid and a very weak base, with low concentration of hydronium and hydroxide ions. Is self ionising. Endothermic self ionising. H2O ⇌ H+ + OH- - obviously most is in molecule form. The product of molar concentrations [H3O+][OH-] is always 1x10-14 at 25˚C. If the H3O+ or OH- increase then the other decreases proportionally.

Adjusting soil acidity

Increased with ammonia nitrate and sulfate fertilisers, decreased with limestone, tested using universal indicator. CaCO3 + 2H3O+ → Ca2+ + 3H2O + CO2

Acid + base → salt + water (neutralisation)

It is still a neutralisation reaction even if not completely neutralised, exothermic. Strong acid + strong base → neutral salt (e.g. NaCl), do not hydrolyse, equilibria so far right, it's completed. Strong acid + weak base → acidic salt (e.g. NH4Cl), hydrolyse producing acidic solutions. Weak acid + strong base → basic salt (NaF), hydrolyse producing alkaline solutions. Hydrolysis - the reaction of a salt with water that results in a change in pH.

pH

Measurement of acidity or H3O+ ions in solution pH = -log10[H3O+], makes writing concentration of H3O+ easier. *Remember that polyprotic calculate as having dissociated completely (no need for steps).

pOH

Measurement of basicity or OH- ions in solution, reverse relationship of pH scale, as pOH increases pH decreases. pOH = -log10[OH-] *at 25˚C pOH + pH of a solution = 14

Weak bases

NH3, C5H5N, (CH3)2NH Weak bases prefer to stay as molecules (they are happy).

Adjusting pool acidity

NaOCl added to pools to kill dangerous microbes, OCl- reacts with water to produce unstable HOCl that kills microbes. This makes pool more alkaline (OH-) so they are neutralised with HCl. Phenol red used to test. OCl- + H2O ⇌ HOCl + OH- H3O+ + OH- → 2H2O

Strong bases

NaOH, KOH, Ca(OH)2 Strong bases readily accept protons, dissociates almost completely, have a tendency to gain protons. Produce OH- easily.

pH metre

Neutral (pH=pOH=7), acidic (pH<7, pOH>7), basic (pH>7, pOH<7) Most acidic have pH slightly less than zero, most basic have pH slightly more than 14. Quantitative, more accurate and reliable than indicators.

Litmus

Red Purple Blue Good to differentiate bases and acids, 5 - 8 is purple, does not change colour in water.

Universal

Red - orange - yellow Green Greenish blue - bluish violet - violet Mixture of indicators, wide range.

Methyl Orange

Red - orangey yellow Yellow Yellow Fairly strong acid, readily losing protons (even when small concentration of hydroxide) Good for differentiating strong and weakly acidic. pH 3.1 - 4.4.

Alkaline solution

Solutions containing hydroxide ions. Water soluble bases are alkalis. E.g. sodium hydrogen carbonate used to neutralise acid burns, anti acid tablets for acid reflux.

Concentration v. Strength

Strength depends on how many of an acid or base are ions (how much it has dissociated), whereas concentration is how much of the ions are there (how much acid or base is dissolved in a given volume of solution). Can measure concentration and strength qualitatively assigning labels, or quantitatively using dissociation constant. E.g. HCl is a strong acid as it readily donates proton, whereas a concentrated solution of HCl can be prepared by bubbling a large amount of hydrogen chloride in a given volume of water.

Acid + metal carbonate → salt + water + carbon dioxide

Strength of reaction depends on strength and concentration of acid. Is exothermic (as proton transfer is exo), use limewater test to check.

Amphiprotic

Substances that donate or accept protons in different contexts (depending on the substance they are reacting with), behaving as either an acid or a base. A subsection of amphoteric with protons. Usually at equilibria, thus can be manipulated by altering temperature and concentration. E.g. HCO3-, H2PO4-, HPO42-, HSO4- E.g. HCl(g) + H2O(l) → Cl-(aq) + H3O+(aq), water gains a proton and acts as a base. NH3(aq) + H2O(l) → NH4+(aq) + OH-(aq), water donates a proton and acts as an acid. I.e. water acts as both depending on solute present, if the solute is a stronger acid than water it will act as a bas, if a stronger base than water it will react as an acid. When an amphiprotic substance is placed in water it acts as both an acid and a base, however one reaction generally dominates identified through measuring the pH.

Strength of conjugate acid-base pairs

The stronger the acid is the weaker the conjugate base. Similarly, the stronger the base is the weaker the conjugate acid.

Conjugate acid-base pairs

Two molecules or ions that differ by one proton (H+).

Applications of acids and bases (examples of neutralisation)

Vinegar used to neutralise box jellyfish stings, HCl to neutralise bases to kill bacteria, sodium hydrogen carbonate used to neutralise acid burns.

Bromothymol blue

Yellow Green Blue Good for small changes in acidity or alkalinity near neutrality. pH range 6 - 7.6.

Polyprotic bases

e.g. CO32- diprotic base.

Diprotic acids

e.g. H2SO4 H2SO4(l) + H2O(l) → HSO4-(aq) + H3O+(aq) - H2SO4 is a strong acid in water as it readily donates a proton, so no H2SO4 molecules are found in solution in this station. HSO4-(aq) + H2O(l) ⇌ SO42-(aq) + H3O+(aq) - HSO4- is a weak acid as it only partially dissociates, with an equilibrium indicating an incomplete reaction. H2CO3 donate two protons.

Indicators

natural (lichens, plant extracts e.g. rose petals, blackberries, cabbage, litmus red purple blue) or synthetic chemical dyes that change colour as the acidity of surroundings change, are organic complex molecules. Are weak acids (represented by HIn), exhibiting one colour when the proton is attached to the molecule, and a different colour when the proton is absent. Can also be bases, the conjugate acid form of the indicator and the conjugate base form another colour. Are examples of equilibrium systems,adding more causing a shift e.g. adding more acid means shift to left, adding more base shift to right. (colour 1) HInd + H2O ⇌ H3O+ + Ind- (colour 2) **Qualitative, the non neutral nature of indicators skews the colours (destructive method), differentiating between colours is also perception based.