NYGH Acids, Bases and Salts Notes

1 Chem!stry Name: …………………… ( ) Class: ……………… Date: …… / …… / …… Essential Notes: Acids, Bases and Salts Names and Formulae of Some Common Acids, Bases and Salts: • Acids: Examples include: hydrochloric acid – HCl, nitric acid – HNO3, sulfuric acid – H2SO4, phosphoric acid – H3PO4, ethanoic acid – CH3COOH • Bases / Alkalis: Note: The term alkali is used to describe a base that is soluble in water, and has been dissolved in water so that it can be used as a solution. Not all bases are soluble in water, and so not all bases are alkalis. Examples include: copper(II) oxide – CuO, sodium hydroxide – NaOH, calcium hydroxide – Ca(OH)2, ammonia – NH3. • Salts: Examples include: sodium chloride – NaCl, copper(II) sulphate – CuSO4, ammonium nitrate – NH4NO3 Definitions: • Acid: An acid is a chemical that dissolves in water to produce hydrogen ions, H+(aq), as the only positive ion: HCl(g) → H+(aq) + Cl–(aq) HNO3(l) → H+(aq) + NO3–(aq) H2SO4(l) → 2H+(aq) + SO42–(aq) Note: Sulfuric acid is said to be dibasic because 1 mol of the acid produces 2 mol of hydrogen ions (also, H2SO4 can donate 2 mol of H+ or 2 mol of H+ can be replaced by a metal to form a salt). The hydrogen ions that are produced do not exist on their own, but bond to water molecules to form hydroxonium ions, H3O+(aq): H+(aq) + H2O(l) → H3O+(aq) Therefore, the complete balanced chemical equation for an acid dissolving in water should be written: HNO3(l) + H2O(l) → H3O+(aq) + NO3 –(aq) (a monobasic acid) CH3COOH(l) + H2O(l) ∏ H3O+(aq) + CH3COO–(aq) (a monobasic acid) H2SO4(l) + 2H2O(l) → 2H3O+(aq) + SO42–(aq) (a dibasic acid) H3PO4(l) + 3H2O(l) → 3H3O+(aq) + PO43–(aq) (a tribasic acid)

2 The dot and cross diagram of the hydroxonium ion is given below: Note: The hydrogen that originated from the acid (shown on the left) is sharing two electrons (two dots) which both belong to the oxygen. This special type of covalent bond is known as a dative covalent bond. An important fact to arise from this is that an acidic chemical will only exhibit its acidic properties when dissolved in water. Therefore, when writing balanced chemical equations, all acids must be in an aqueous state, (aq). • Base / Alkali: Bases tend to be metal oxides and metal hydroxides. Ammonia, NH3, is also basic. As opposed to producing hydrogen ions, which is the property of an acid, a base will accept hydrogen ions to form water as one of the products: CuO(s) + 2H+(aq) → Cu2+(aq) + H2O(l) NaOH(s) + H+(aq) → Na+(aq) + H2O(l) A base which is soluble in water is described as an alkali. Examples of alkalis include sodium hydroxide, NaOH, and ammonia, NH3. Alkalis tend to dissolve in water to produce hydroxide ions, OH –(aq): NaOH(s) → Na+(aq) + OH–(aq) NH3(g) + H2O(l) → NH4+(aq) + OH–(aq) Hydroxide ions (from alkalis) react with hydrogen ions (from acids) to form water. This is known as a neutralisation reaction. The ionic equation is: OH–(aq) + H+(aq) → H2O(l) • Salt: Salts are ionic compounds, usually composed of a positive metal ion bonded to a negative non-metal ion. An exception would be a salt that contains a positive ammonium ion, NH4+, in place of the positive metal ion. Salts are often prepared by replacing the hydrogen ion(s) of an acid with a metal ion(s): sodium chloride: NaCl composed of Na+ and Cl– copper(II) sulphate: CuSO4 composed of Cu2+ and SO42– ammonium nitrate: NH4NO3 composed of NH4+ and NO3–

3 Properties of Acids: • Acids turn blue litmus paper red. • Acids turn universal indicator red / orange / yellow (depending upon the pH value of the acid). • Acids have pH values less than 7. A pH of exactly 7 is neutral. pH = –log10[H+] Where [H+] is the hydrogen ion concentration, mol/dm3 Example of a pH calculation: If [H+] = 0.001, then pH = –log10 0.001 pH = –(–3) pH = 3.00 (3 s.f.) • An aqueous solution of an acid will conduct electricity because it contains mobile ions (charge carrying particles) which are free to move to the electrode of opposite charge. For example, an aqueous solution of sulfuric acid contains a mixture of mobile hydrogen ions, H+, which will be attracted to the negative electrode (cathode) and mobile sulphate ions, SO42-, which will be attracted to the positive electrode (anode). • Acids react with metals to produce a salt and hydrogen as the products: acid + metal → salt + hydrogen nitric acid + magnesium → magnesium nitrate + hydrogen 2HNO3(aq) + Mg(s) → Mg(NO3)2(aq) + H2(g) Ionic equation: 2H+(aq) + Mg(s) → H2(g) + Mg2+(aq) This is a redox reaction because the metal is oxidised while the hydrogen is reduced. This reaction can also be considered as a displacement reaction. The hydrogen of the acid will only be displaced by a metal that is more reactive than hydrogen in the reactivity series, i.e. copper, gold and silver will not displace hydrogen from an acid (there will be no observed reaction when copper, gold or silver are added to an acid). Note: Hydrogen gas extinguishes a burning splint with a “pop” sound. • Acids react with carbonates to produce a salt, water and carbon dioxide as the products: acid + carbonate → salt + water + carbon dioxide sulfuric acid + sodium carbonate → sodium sulphate + water + carbon dioxide H2SO4(aq) + Na2CO3(aq) → Na2SO4(aq) + H2O(l) + CO2(g) Ionic equation: 2H+(aq) + CO32–(aq) → H2O(l) + CO2(g)

4 The salt is formed when the hydrogen of the acid is replaced by the metal of the carbonate. This will work for any metal carbonate, including copper(II) carbonate and silver carbonate. Note: Carbon dioxide gas forms a white precipitate when bubbled through lime water. • Acids react with bases / alkalis to produce a salt and water as the products: acid + base → salt + water ethanoic acid + calcium hydroxide → calcium ethanoate + water 2CH3COOH(aq) + Ca(OH)2(aq) → (CH3COO)2Ca(aq) + 2H2O(l) Ionic equation: H+(aq) + OH–(aq) → H2O(l) The salt is formed when the hydrogen of the acid is replaced by the metal of the metal oxide / hydroxide. This will work for any metal oxide / hydroxide, including copper(II) oxide and silver hydroxide. Note: When a dibasic or tribasic acid reacts, it is possible for the hydrogen ions of the acid to be replaced by metal ions one-at-a-time, i.e. the hydrogen ions of a dibasic or tribasic acid do not need to be replaced all-at-once. • Example, phosphoric acid reacting with sodium hydroxide: • Replace first H+ with Na+ to form monosodium dihydrogen phosphate (an acid salt): H3PO4(aq) + NaOH(aq) → NaH2PO4(aq) + H2O(l) • Replace second H+ with Na+ to form disodium monohydrogen phosphate (an acid salt): NaH2PO4(aq) + NaOH(aq) → Na2HPO4(aq) + H2O(l) • Replace third H+ with Na+ to form trisodium phosphate (or just sodium phosphate): Na2HPO4(aq) + NaOH(aq) → Na3PO4(aq) + H2O(l) • Overall: H3PO4(aq) + 3NaOH(aq) → Na3PO4(aq) + 3H2O(l) • Acids fall into two categories, strong acids and weak acids. Strong acids include hydrochloric acid, nitric acid and sulfuric acid. These acids fully ionize or dissociate when dissolved in water. For example, if 100 molecules of nitric acid, HNO3, are added to water, all of them will ionize to from 100 hydrogen ions, H+, and 100 nitrate ions, NO3–. There will be no nitric acid molecules left intact: HNO3(l) → H+(aq) + NO3 –(aq)

5 Weak acids tend to be organic acids (ones that contain carbon) such as ethanoic acid and citric acid. These acids only partially ionize or dissociate when dissolved in water. For example, if 100 molecules of ethanoic acid, CH3COOH, are added to water, only 5 of them will ionize to from 5 hydrogen, H+, ions and 5 ethanoate ions, CH3COO–. The remaining 95 ethanoic acid molecules will remain intact: CH3COOH(l) Ý H+(aq) + CH3COO–(aq) The Ý symbol indicates that th