Acids and Bases Part 1 Properties of Acids

Acids and Bases Part 1

Properties of Acids 1. 2. 3. Aqueous solutions of acids have a sour taste. Acids change the color of acidbase indicators. Some acids react with active metals to release hydrogen gas, H 2.

Properties of Acids (cont. ) 4. Acids react with bases to produce salts and water. 5. Some acids conduct electrical current.

Binary Acids l An acid that contains only 2 different elements: hydrogen and one of the more electronegative elements. l Examples: – HF – HCl – HBr

Oxyacids l An acid that is a compound of hydrogen, oxygen, and a third element, usually a nonmetal. – HNO 3, nitric acid l. Name of a ternary acid is based on this anion.

Acids l Name the following acids: – HBr. O -HF – HBr. O 3 -H 2 SO 4 l Write formulas for the following acids: – Nitrous -Phosphoric – Hydrochloric -Acetic

Common Industrial Acids l Sulfuric acid-metal production l Nitric acid-explosives l Phosphoric acid-flavor drinks l Hydrochloric acid-cleaning l Acetic acid-plastics

Properties of Bases 1. 2. 3. Aqueous solutions of bases taste bitter. Bases change the color of acidbase indicators. Dilute aqueous solutions of bases feel slippery.

Properties of Bases (cont. ) 4. Bases react with acids to produce salts and water. 5. Bases conduct electrical current.

Arrhenius Acids l Arrhenius theorized that acids and bases must produce ions in solution. – Arrhenius acid a chemical compound that increases the concentration of hydrogen ions, H+, in aqueous solution.

Hydronium Ion l The H+ ion attracts other molecules so strongly that it does not exist alone. l It typically becomes covalently bonded to oxygen to form H 3 O+, the hydronium ion. H 2 O(l) + HCl(l) H 3 O+(aq) + Cl-(aq)

Arrhenius base l Arrhenius base a substance that increases the concentration of hydroxide ions, OH-, in aqueous solution.

Examples l Arrhenius Acids HCl(l) + H 2 O(l) H 3 O+(aq) + Cl-(aq) l Arrhenius Bases Na. OH(s) Na+(aq) + OH-(aq) l Which one is an acid, which one is a base? HNO 3(aq) + KOH(aq) KNO 3(aq) + H 2 O(l)

Strength of Acids l Depends on the extent to which the acid dissociates. l Strength of an acid depends on the polarity of the bond between hydrogen and the element to which it is bonded and the ease with which that bond can be broken.

Strong Acids l Strong acid one that ionizes completely in aqueous solution. – HCl. O 4 perchloric acid – HCl hydrochloric acid – HNO 3 nitric acid

Weak Acids l Weak acid: Acids that are weak electrolytes – H 3 PO 4 + H 2 O <=> H 3 O+ + H 2 PO 4 - – H 2 CO 3 + H 2 O <=> H 3 O+ + HCO 3 -

Acid-Base Theories l Bronsted-Lowry acids and bases l Lewis acids and bases

Bronsted-Lowry Theory l Bronsted-Lowry acid a molecule or ion that is a proton donor – All acids defined by Arrhenius are Bronsted-Lowry acids as well.

Bronsted-Lowry Base l Bronsted-Lowry base a molecule or ion that is a proton acceptor

l In a Bronsted-Lowry acid-base reaction, protons are transferred from one reactant (the acid) to another (the base). l Example: HCl + NH 3 NH 4+ + Clacid base

Examples Identify each compound as an acid or a base. HCl + H 2 O H 3 O+ + Cl. H 2 O + NH 3 NH 4+ + OHH 2 CO 3 + H 2 O HCO 3 - + H 3 O+

l Monoprotic acid an acid that can donate only one proton (hydrogen ion) per molecule. – HCl. O 4; HCl; HNO 3 l Polyprotic acid an acid that can donate more than one proton per molecule. – H 2 SO 4; H 3 PO 4

Acids and Bases Part II

Lewis Theory l Lewis acid atom, ion, or molecule that accepts an electron pair to form a covalent bond. l Lewis base atom, ion, or molecule that donates an electron pair to form a covalent bond.

l. A Lewis acid-base reaction is the formation of one or more covalent bonds between an electron-pair donor and an electron-pair acceptor.

Examples H 3 N + BF 3 H 3 NBF 3 + F- => BF 4 -

Type Acid Base Arrhenius H+ or H 30+ producer Proton (H+) donor Electron pair acceptor OHproducer Proton acceptor Electron pair donor Bronsted. Lowry Lewis

So which theory do we use? l We will focus mainly on the Bronsted -Lowry theory of acids and bases. – But you still need to be able to define and recognize acids and bases according to the other two theories!!!

Conjugate Bases l. A conjugate base is the species that remains after a Bronsted-Lowry acid has given up its proton. – HF(aq) + H 2 O(l) F-(aq) + H 3 O+(aq)

Conjugate Acids l. A conjugate acid is the species that is formed when a Bronsted-Lowry base gains a proton. – HF(aq) + H 2 O(l) F-(aq) + H 3 O+(aq)

Practice Problems l For each reaction listed, identify the proton donor or acid and the proton acceptor or base. Label each conjugate acid-base pair. – HCO 3 -(aq) + H 2 O(l) H 2 CO 3(aq) + OH-(aq) – HNO 3 + SO 42 - HSO 4 - + NO 3 -

Strength of Conjugate Acids and Bases l The extent of a reaction between a Bronsted-Lowry acid and base depends on the relative strengths of the acids and bases involved. l The stronger an acid is, the weaker its conjugate base; the stronger a base is, the weaker its conjugate acid.

Amphoteric Compounds l Any species that can react as either an acid or a base is described as amphoteric. l Water acts as a base in the following reaction: – H 2 SO 4(aq) + H 2 O(l) H 3 O+(aq) + HSO 4 -(aq)

l Water acts as an acid in the following reaction: – NH 3(g) + H 2 O(l) NH 4+(aq) + OH-(aq) l Amphoteric substances act as either an acid or base depending on the strength of the acid or base with which they are reacting.

Neutralization Reactions l Neutralization Strong Acid + Strong Base H 2 O + salt A salt is an ionic compound composed of a cation from a base and an anion from an acid.

Practice Problems l Complete the following neutralization reactions and balance them. – HCl(aq) + Na. OH(aq) – HNO 3(aq) + KOH(aq) – Ca(OH)2(aq) + HNO 3(aq) – Mg(OH)2(aq) + HCl(aq)

l l l What is the relationship between the strength of an acid and its conjugate base? What is the relationship between the strength of a base and its conjugate acid? What is meant by the term amphoteric?