Chapter 5 and 17 Acids and Bases Introduction

General Rule: 1. If the oxide is covalent and a strong bond holds the oxygen – acidic solutions are produced Ex. SO 3 + H 2 O  H 2 SO 4 2. If the oxide is ionic – the compound will produce a basic solution in water. Ex. CaO + H 2 O  Ca(OH) 2 What will make an acid/base?

Properties of Acids: Sour taste Change color of indicators Some react with metals to produce H 2 gas Are neutralized by the reaction with a base Some conduct electricity

2 factors that determine the strength of an acid: 1. Binary acids - The strength of a bond– the stronger the bond, the weaker the acid (harder to dissociate) 2. Oxyacids - The polarity of the bond – the more oxygens – the more polar the molecule – stronger the acid The more + charge of the metal cation in a coordination compound – the stronger the acid – increased polarity The more electronegative metal in an oxyacid – stronger the acid due to increased polarity

Common Uses for Acids: A. Sulfuric acid – most commonly used – in making of metals, paper, paints Attracts water – dehydration agent B. Nitric – rarely used – very unstable – has a suffocating odor, stains skin, burns Used to make explosives, rubber, plastics

C. Phosphoric – used in fertilzers, detergents, ceramics, diluted in pop D. Hydrochloric – digestion, cleaning agent, acidity in pools E. Acetic – (glacial acetic acid – concentrated - will freeze at 17C) Vinegar is 4-8% acetic acid Used in plastics and food supplements

Types of Acids: Monoprotic – have one acidic H + - ex. HCl Diprotic – have 2 acidic H + - ex. H 2 SO 4 Triprotic – have 3 acidic H + - ex. H 3 PO 4 Polyprotic – acids that can donate more than 1 acidic H + Organic – have a carbon backbone – usually very weak – have only 1 acidic hydgrogen Hydrohalic – acidic proton is attached to a halogen – Ex. HCl or HF

Properties of bases: Bitter taste Change colors of acid/base indicators Feel slippery Are neutralized by the reaction of an acid – produce salt and water Electrolytes Neutralization – when a strong acid and base react they neutralize each other to form a salt (ionic compound) and water

1. Arrhenius concept – acids produce H + in aqueous solutions and bases produce OH - Only applies to acids in aq solutions and bases that contain OH - 3 ways to define an acid/base

2. Bronsted-Lowry Model Acid is a proton donor Base is a proton acceptor Hydronium ion – H 3 O + Polyprotic acids only dissociate one acid at a time.

General Bronsted Lowry Reaction: HA(acid) + H 2 O (base)  H 3 O + (Conjugate acid) + A - (conjugate base) Conjugate base – everything that remains of the acid after the proton is lost – will have a neg. charge Conjugate acid – formed when the proton is transferred to the base – (will have a + charge) Conjugate acid/base pair – 2 substances that are related due the accepting/donating of a proton. HA and A - (acid and its conjugate base) and H 2 O and H 3 O + (the base and its conjugate acid)

The stronger the acid; the weaker the conjugate base. The stronger the base, the weaker the conjugate acid. Amphoteric (amphiprotic) – can act like an acid or a base Ex. Water Autoionization – transfer of a proton from one molecule to another of the same substance to produce an acid and a base

3. Lewis Concept Lewis acid – electron pair acceptor (does not have to be H) Lewis base – electron pair donor (does not have to be H) Will form 1 product – acid –base adduct Look for bases that are anions or neutral molecules that have lone pairs Look for acids that are cation or neutral molcules with empty valence orbitals such as B and Be

Acid-Base Indicators Compounds whose color changes when the pH changes These are weak acids/bases Will be their original color in acidic solution and a different color in a basic solution as the indicator dissociates Universal indicators – have several different indicators mixed together – will show different colors at different pHs – fairly accurate

pH meters Used if the exact pH is needed – measures the voltage between 2 electrodes placed in the solution The voltage changes as the H + concentration changes

Titrations Used to determine the concentration of an unknown acid/base by a known acid/base Equivalence point (Stoichiometric point) when the concentrations of the unknown acid/base and the known acid/base are equal – determined with an indicator or pH meter Endpoint – point during a titration where an indicator changes color A good indicator’s endpoint matches the equivalence point of the titration

How to determine the equivalence point range: 1. Strong Acid with a Strong Base – pH will be 7.00 at this point – neutral 2. Weak acids with a strong base – pH will be greater than 7 C. Weak bases with a strong acid – pH will be less than 7 D. Weak acid with a weak base - beyond the scope of this class

pH Curve (Titration Curve): Plot of the pH of the solution as a function of the amount of titrant added. Can use millimol (mmol) per milliliter to describe titrations since the quantities are usually small and burets are in mL Molarity = mmol/mL

2 important facts about titration curves: 1. It is the AMOUNT of the acid, not the strength that determines the amount of base needed to reach the equivalence point. 2. The pH value at the equivalence point IS affected by the acid strength. The weaker the acid, the greater the pH at the equivalence point.

Standard solution – the known solution Primary standard – the highly purified solid used to check the concentration of the known solution

Steps on how to determine the concentration of an unknown through titration: 1. Write the balanced neutralization reaction. 2. Determine the moles of the known acid/base 3. Determine the moles of unknown used during the titration. 4. Determine the molarity of the unknonwn.

Equilibrium Constant – Ka and Kb Strong acid and bases – equilibrium lies far to the right – completely dissociaties at equilibrium will make a weak conjugate base/acid – water is the main proton acceptor if an acid or proton donor if a base. Large Ka if an acid or Kb if a base. K>1 Weak acid or base – equilibrium lies to the left – will hardly dissociate at equilibrium. conjugate base or acid is very strong – conjugate base is the main proton acceptor or conjugate acid is the main proton donor. Small Ka if weak acid or Kb if a weak base

**Stronger the acid – the weaker its conjugate base is** There is a competition taking place for the H + between water and the conjugate base. If water is stronger – equilibrium lies far to the right. If the conjugate base is stronger – equilibrium lies to the left.

Acid – Base Properties of Salts Salt – ionic compound – will break into ions when they dissociate in water Salts that have cations of strong base (Na + ) and anions of strong acids (Cl - ) have no effect on the H + concentration – therefore, they are neutral – pH = 7.00

Salts of weak acids The conjugate base of a weak acid has an affinity for protons – therefore conjugate base affects the pH. A basic solution is formed if the anion of the salt is the conjugate base of a weak acid. Anions from polyprotic acids can act as an acid or a base.

An acidic solution will be formed if the anion is NOT a base and the cation is the conjugate acid of a weak base – usually only ammonium and its derivitatives If a salt contains a charged metal – will from a complex ion Ex. Al makes Al(H 2 O) 6 +3 – this is a conjugate acid Basic if written as [Al(H 2 O) 5 (OH)] +2 Salts of Weak Bases

If both ions of the salt are from weak acids/bases Just compare the K values 1. If K a > K b – acidic 2. If K a < K b – basic 3. If K a = K b – neutral

K a * K b = K w Works for a weak acid and its conjugate base K a – weak acid; K b is the conjugate base pK a = -logK a

Predicting the direction: The reaction will always move from the stronger acid/base to the weaker acid/base If a weak acid and a weak base – must compare the Ka and Kb values of the conjugate acid and base.