Chapter 5 and 17

1. Chapter 5 and 17 Acids and Bases Introduction

2. What will make an acid/base? • General Rule: • 1. If the oxide is covalent and a strong bond holds the oxygen – acidic solutions are produced • Ex. SO3 + H2O  H2SO4 • 2. If the oxide is ionic – the compound will produce a basic solution in water. • Ex. CaO + H2O  Ca(OH)2

3. Properties of Acids: • Sour taste • Change color of indicators • Some react with metals to produce H2 gas • Are neutralized by the reaction with a base • Some conduct electricity

4. 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

5. 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

6. 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

7. Types of Acids: • Monoprotic – have one acidic H+ - ex. HCl • Diprotic – have 2 acidic H+ - ex. H2SO4 • Triprotic – have 3 acidic H+ - ex. H3PO4 • 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

8. 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

9. 3 ways to define an acid/base • 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-

10. 2. Bronsted-Lowry Model • Acid is a proton donor • Base is a proton acceptor • Hydronium ion – H3O+ • Polyprotic acids only dissociate one acid at a time.

11. General Bronsted Lowry Reaction: • HA(acid) + H2O (base)  H3O+ (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 H2O and H3O+ (the base and its conjugate acid)

12. 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

13. 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

14. 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

15. 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

16. 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

17. 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

18. 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

20. 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.

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

22. 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.

23. 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

24. **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.

25. 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

26. 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.

27. Salts of Weak Bases • 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(H2O)6+3 – this is a conjugate acid • Basic if written as [Al(H2O)5(OH)]+2

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

29. Ka * Kb = Kw • Works for a weak acid and its conjugate base • Ka – weak acid; Kb is the conjugate base • pKa = -logKa

30. 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.