Acids bases & salts

1. Acids bases & salts

2. Objectives • State the Bronsted-Lowry definition of acids and bases • Identify the common physical and chemical properties of acids and bases • Explain what dissociation constants indicate about an acid or base • Use experimental data to calculate a dissociation constant

3. Properties of acids/bases • Taste – acid comes from latin meaning sour or tart/bases are bitter – soap • Touch most dilute acids feel like water although they sting on broken skin bases feel smooth, soothing and slippery except in your eyes (soap) • Reactions with metal acids react bases do not react

4. Properties continued • Electrical conductivity water is poor – HCl is good, NaOH is good both are electrolytes • Indicators turn color – acid turns litmus paper from blue to red base turns from red to blue • Neutralization reaction between an acid and a base get salt and water (double replacement)

5. Arrhenius definition • An acid is a substance that dissociates in water to produce hydrogen ions. • A base is a substance that dissociates in water to produce hydroxide ions • A salt is an ionic compound formed from any cation other that H+ and any anion other than OH- or O-2

6. Arrhenius continued • Acids react with metals to produce H2 gas • Mg +2H+ -> Mg+2 + H2 • oxidation reduction reaction

7. Bronsted – Lowry definitions • An acid is any substance that can donate H+ ions • A base is any substance that can accept H+ ions a Bronsted-Lowry acid is a proton donor and a base is a proton acceptor

8. Hydronium Ion • H+ is strongly attracted to the electrons of surrounding water molecules • H+ + H2O -> H3O+ • More correct HCl + H2O -> H3O++ Cl- • In this case HCl is the Bronsted-Lowry acid and water is the base • We still describe a solution of HCl as acidic!

9. NH3 + H2O -> NH4+ + OH-Ammonia is the H+ acceptor water is the H+ donor (acid) Amphoteric a substance that can act as either an acid or a base

10. Conjugate Acid-Base Pairs • NH3 + H2O  NH4+ + OH- • In 1 direction water is the acid in the reverse reaction it is the base. • These cmpds become conjugate acids and conjugate bases when HCl loses an H+ ion to become its conjugate base Cl- when the conjugate base of water is the hydroxide ion OH- • When ammonia gains H+ to become its conjugate acid NH4+ and the conjugate acid of OH- is H2O

11. Conjugate pairs • A pair of compounds that differ by only one H+ ion such as H2O and OH- or NH3 and NH4+ are called conjugate acid base pairs • NH3 + H2O  NH4+ + OH- • Base acid conj. Acid conj. base

12. Determining the strengths of acids and bases • A strong acid HCl readily transfers hydrogen ions to water to form H3O+ • If you place 1M of HCl in 1 liter of water you would form 1 M H3O+ ions and 1 M Cl- ions

13. weak acids • A weak acid does not readily transfer H+ ions • 1 mole of acetic acid in 1 liter of water only .4% of the acetic acid molecules would form H3O+ and C2H3O2-. Which means that 99.6% of the acetic acid molecules do not dissociate.

14. To show a strong acid from a weak acid use arrows • HCl + H2O  H3O+ + Cl- • HC2H3O2 + H2O <-> H3O+ + C2H3O2-

15. Strong and Weak Bases • The most widely used commercial base is CaO. When CaO is dissolved in water the O-2 ions react completely with H2O to form OH- ions. • O-2 + H2O  2OH- use a single arrow

16. Strength of conjugate acid – base pairs • The stronger the acid the weaker its conjugate base. • The stronger the base the weaker its conjugate acid.

17. The acid dissociation constant Ka • Weak acid HA • HA + H2O <--> H3O+ + A- • Keq = • For a 1 M solution of a typical weak acid may be only .007% of the water molecules react. Move the water to the left side of the equation • The higher the Ka the more the reaction goes to the right . The greater the Ka the stronger the acid • Weak acids have a Ka less than 1

18. Diprotic acids • 2 step dissociation

19. Base dissociation constant Kb • The base dissociation constant is a measure of the strength of a base • Do calculations of dissociation constants

20. Objectives • Explain what most acidic hydrogen atoms have in common • Explain what most bases have in common • Describe how acids are named

21. Naming and identifying acids and bases • 1st all H’s are not acidic CH4 • As a rule an acidic hydrogen already has a slight positive charge while is it part of a molecule. (It is in the positive side of a polar covalent bond) • Usually bonded with O, N, or a halogen

22. 3 types of acids • Binary acids - H and 1 other element usually 6A or 7A • Strong HCl, HBr, and HI • Weak HF, H2S and H2Se • Oxy acids contain H, O and 1 other element • H2SO4, HNO3 and H3PO4 • Carboxylic Acid – organic acids COOH group • Acetic acid HC2H3O2 vinegar

23. Bases • A Bronsted-Lowry base always contains an unshared pair of electrons NH3 attracts H+ • Anions: remember conjugate bases HCl Cl- • Weak Cl- Br- I- NO3- HSO4- CIO4- • Strong O-2 OH- PO4-3 and CO3-2 • Amines N has an unshared pr. Of electrons

24. Naming Acids and Bases • If the name of an anion ends in ide the name of the acid that produces it includes the name of the anion, a hydo prefix • Hydrochloric acid – all binary acids • If the name of an anion ends in ate use and ic ending Nitric acid, carboxylic acid • Anion ends in ite SO3-2 sulfite ion H2SO3 sulfurous acid