Chemistry 141 Wednesday, October 4, 2017 Lecture 13 Solution Chemistry, Part 3: Acids and Bases, Oxidation States

Objectives for today Begin to gain chemical intuition about reactions in solution Identify and differentiate between strong and weak acids and bases Measure acid strength using pH Assign oxidation states to elements in compounds Investigate reactions in which electrons are transferred (redox reactions)

Types of solution reactions Precipitation reactions make a solid product, almost always an insoluble salt. Acid–base reactions require both an acid and a base as reactants usually form water as a product Redox reactions characterized by changes in the oxidation states of elements

Acids An acid is a proton (H+) donor – it dissociates to give H+ when dissolved in water Strong acids totally dissociate in water – have low pH Weak acids partially dissociate in water – have pH < 7 Some acids (like H2SO4) are polyprotic – they contain more than one acidic H in general, the first H is strongly acidic, and other H’s are weakly acidic

Bases A base is a proton (H+) acceptor Strong bases totally dissociate in water – have high pH Some weak bases do not contain OH, but react with water to produce some OH in solution – have pH > 7 sodium carbonate Na2CO3  2 Na+ + CO32– CO32– + H2O  HCO3– + OH– ammonia NH3 + H2O  NH4+ + OH–

Acids and Bases: pH pH is a measure of the H+ ion concentration in a solution pH = -log10[H+] [H+] = 10-(pH) where [H+] is the H+ ion concentration in moles/liter (M) The analogous property for OH– is pOH = -log[OH–] Water is a covalent molecule, hence in liquid water, a very few water molecules dissociate to give H+ + OH– H2O  H+ + OH– The pH of pure water is 7 – a solution with this pH is neutral [H+] = [OH–] =10-7 pH+pOH = 7 + 7 = 14 For any solution (not just neutral solutions), pH + pOH always equals 14 there are 10-7 moles H+ per liter of water

Rules for determining oxidation states The oxidation state of an atom in its elemental form is 0. The oxidation state of a monatomic (free) ion is equal to its charge. The sum of the oxidation numbers of the atoms in any uncharged compound is 0. The sum of the oxidation numbers of the atoms in a charged species (such as a polyatomic ion) is equal to the charge of the species. Within compounds, the following rules apply in order : Alkali metals have oxidation number +1 (e.g., NaCl). Alkaline earth metals have oxidation number +2 (e.g., BaCl2). Hydrogen (H) has oxidation number +1, except in compounds with alkali metals or alkaline earth metals. Fluorine (F) has oxidation number –1. Oxygen (O) has oxidation number –2, except in compounds with fluorine. The other halogens have oxidation number –1, except in compounds with fluorine or oxygen.

Oxidation – Reduction (Redox) Reactions Oxidation – Reduction reactions involve a change in oxidation states Two elements must change oxidation states electrons are transferred total number of electrons must be conserved If one element is oxidized, another must be reduced Oxidation – increase in oxidation state, loses electron(s) Reduction – decrease in oxidation state, gains electron(s) The oxidizing agent is reduced, and reducing agent is oxidized in the reaction

Displacement Reactions

Activity Series Elements higher on the activity series are more reactive. They are more likely to exist as ions.