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Acids and Bases

Learn about the properties of acids and bases, including pH range, taste, reactions, and common examples. Understand the Arrhenius and Bronsted-Lowry definitions, as well as Lewis acids and bases.

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Acids and Bases

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  1. Acids and Bases

  2. Properties of Acids • They are corrosive • pH range is 0 to 7, lower the pH, more acidic the acid • Taste sour • Turns litmus paper red • Often reacts with metals to produce H2 gas • Often produces H+ in water. Example; HCl  H+ + Cl-

  3. Symbol used for acids

  4. Properties of Bases • pH range is from 7 to 14, higher the pH more alkaline the base • They are caustic • Taste bitter • Turns litmus paper blue • Feels slippery • Group 1 and 2 metals make common bases. Example: Na + H2O  NaOH + 1/2H2

  5. The pH of some common acid and bases

  6. Common Acids • Binary or Halide Acids • HCl, hydrochloric acid: is stomach and pool acid • HF, hydrofluoric acid: used to etch glass

  7. Common acids continued • Oxyacids: acids with oxygen (more oxygen- stronger the acid) • HNO3, Nitric acid, used for making explosives, when it reacts with metals it produces NO2gas instead of H2 • H2SO4, sulfuric acid: used in car batteries, used in many industrial process as a dehydrator.

  8. Chemical burn from H2SO4 • AHHHHHH SATAN!!!!!!! =]

  9. Reaction of Cu with HNO3

  10. Oxyacids continued • H2CO3, carbonic acid, found in carbonated sodas, and in rainwater. It is responsible for the sour taste, or sting in sodas. It is also found in blood to a very small extent. • H3PO4, phosphoric acid, used as lime scale remover, and as flavoring in the colorless sodas like Sprite.

  11. Acids continued • Organic acids • Carbon based acids. • Structure ends in COOH • COOH is called a carboxyl group. • CH3COOH, ethanoic acid or vinegar. Used as a preservative and flavoring O C-OH

  12. Bases • Group 1 bases • Strongest bases • NaOH, sodium hydroxide, and KOH, potassium hydroxide, both of these bases are used to make commercial cleansers and drain cleaners. • Group 2 bases • Weak bases, • Ca(OH)2 calcium hydroxide, make limewater which is used to test for CO2,,used to make antacids.

  13. Bases • Amines. • This group has a -NH2 at the end of a molecule • These are formed from decomposing proteins • They have a “fishy” smell • NH3, ammonia, this is used as glass cleaner and to make fertilizers. • Examples CH3NH2, methyl amine, CH3CH2NH2 ethyl amine. Etc.

  14. Bases • Carbonates (they are antacids) • Have CO3-2 in them • This group produces CO2 when they react with acids. • NaHCO3 sodium bicarbonate, or baking soda • CaCO3, calcium carbonate, or baking powder. • MgCO3 and Al2(CO3)3 are used as antacids, for example: MgCO3 + 2HCl MgCl2(aq) + H2O(l) + CO2(g)

  15. Arrhenius definition of acids and bases • Acids produce a hydrogen ion(H+) or hydronium ion (H3O+) when they dissolve in water. (An H+ ion is considered too reactive to exist so an H3O+ ion is used) • Acids dissolving in water producing an H+ • HCl  H+ + Cl- • HNO3  H+ + NO3- • H2SO4  2H+ + SO4-2 • CH3COOH  CH3COO- + H+

  16. Carbonates are in antacids Turner's stash

  17. Arrhenius definition continued • Acids dissolving in water form a hydronium ion: • HCl + H2O  H3O+ + Cl- • HNO3 + H2O  H3O+ + NO3- • H2SO4 + 2H2O  2H3O+ + SO4-2 • CH3COOH + H2O  CH3COO- + H3O+ • Note: H2O is written in the reaction when H3O+ is used, but not when H+ is used.

  18. Arrhenius definition of bases • Bases dissolve in water to produce an OH- called a hydroxide ion. • NaOH  Na+ + OH- • Ca(OH)2  Ca+2 + 2OH- • NH3 + H2O  NH4+ + OH- • Note: H2O is written in the reaction only with amines

  19. Bronsted-Lowry def. of Acids/Bases • Acids are proton donors • Bases are proton acceptors • Ex: HNO2 + H2O H3O+ +NO2- • Ex: NH3+ HCO3-   NH4+ + CO3-2 • Ex: HCO3- + HSO4-   SO4-2 + H2CO3

  20. Determining the conj. Acid of a base • --Add a H+ • Base Conj. Acid • HCO3- H2CO3 • H2O H3O+ • OH- H2O

  21. Determining the conjugate base of an acid. • --remove an H+ • Acid Conj. Base • H2O OH- • HSO4- SO4-2 • NH4+ NH3

  22. Continued Bronsted-Lowry • Stronger the acid/ base, the weaker its conjugate base/ acid. • Amphoteric or Amphiprotic: A substance which can be either an acid or a base. • Example is H2O. • Water as an acid: H2O + Cl-  OH- + HCl • Water as a base: H2O + HCl  H3O+ + Cl-

  23. Conjugate acid-base pairs

  24. Definition of a Lewis acid and base • Acids are electron pair acceptors in a dative covalent bond. • Bases are electron pair donors in a dative covalent bond. • These acids are reserved for those molecules/ions which make dative covalent bonds but are NOT ALREADY A BRONSTED-LOWRY ACID. • Examples: NH3 is a base because the nitrogen has a pair of electrons which can be used in a dative bond. AlF3 can be a Lewis acid since Al has a vacant pair of orbitals that can accept e-.

  25. Lewis acids will include substances that are not considered typical acids. • Examples: • Some metal ions. Ex Group 3, Trans. metals

  26. Strong acids and bases completely dissociate in water

  27. Strong and weak acids and bases • Strong acid/base when it dissolves in H2O, nearly 100% dissociates into ions. • Ex: HCl H+ + Cl- nearly 100% of the molecules break up into ions • Ex: NaOH Na+ + OH- nearly 100% of the molecules break up (Remember: the H+ and OH- gives the properties associated with Acids and Bases.) ▪ For weak acid and bases, when they dissolve in water only a small % will dissociate to form ions. CH3COOH(aq) CH3COO-(aq)+ H+(aq) Very little of the acid forms ions. Same condition for weak bases. ▪ Both the strength and the concentration of the acid or base determines its pH and its harmfulness. So when working with acids and bases one needs to be concerned with 1. Is the acid or base strong or weak? 2. Is the acid or base concentrated or dilute?

  28. WWCND • How burgers are made

  29. The pH equation • The pH equation is: pH = -log [H+] • Where [H+] is the molar concentration.

  30. The pH scale • Development of the pH scale • Based upon the dissociation of water: H2O  H+ + OH- • Based upon the equilibrium of water: Kw = [H+] [OH-] • at equilibrium the concentrations of H+ and OH- are each 1x10-7M (this was experimentally determined) • Substitute these concentrations into the Kw expression: Kw = [1x10-7][1x10-7]; this = 1x10-14

  31. Development of the pH scale cont. • According to the rules of equilibrium changing conc does not change the Kw constant. So if acid is added to water, the H+ conc goes up, and the OH- goes down, but K equals 1x10-14. The same thing happens if a base is added to water. Knowing that Kw is always 1x10-14, then • The largest conc for either an H+ or OH – is 1M • The smallest conc H+ or OH – is 1x10-14 that we work with • the –log of 1M = 0 the low end of the pH scale • the –log of 1x10-14 = 14 the high end of the pH scale • Also, pure H2O has a conc of H+ = 1x10-7, the –log = 7, pure H2O has a pH of 7 or it is neutral

  32. Acid and Base Neutralizations • General equation: acid + base  salt+ H2O + energy • Ex: HCl+ NaOH  H2O + NaCl • CH3COOH + NaOH HOH + NaCH3COO • H2 SO4 + 2NaOH  2 HOH + Na2SO4 • HNO3 + NaOH  NaNO3 + H2O • 2HCl+ Ca(OH)2  CaCl2 + 2H2O

  33. Titration Curves for strong acid and bases

  34. Titration curve for strong acid and bases

  35. Explanation of a titration curve • A titration involves an acid base reaction: • Example NaOH + HCl  NaCl + NaOH • All titration curves have the same shape but they do not all start at the same pH • Why this shape? • Start with 1M HCl titrated with 1M NaOH • Initial pH=O, titrate 10% of 1MHCl, 90% of HCl remains. [H+]= .9M or pH= .05 • Titrate 50% of 1M HCl, 50% remains [H+]= .5M or pH= .3 (not a great change in pH yet!) • Titrate 90% of 1MHCl , 10% remains, [H+]=.1M or pH=1 • Titrate 100% of HCl, acid and bases are now equal, pH=7 the cure shoots up until excess NaOH is added. Now curve gradually increases again with the addition of more NaOH

  36. Tiration curves with weak acids and bases.

  37. IB optional material: Equilibrium and weak acids and bases • Since weak acids and bases disassociate partially: • They are reversible and have a measurable equilibrium. • Their pH cannot be based upon the initial molarity of the acid or base. An equilibrium equation must be used to measure it.

  38. The Ka or Kb (equilibrium constant) of weak acids or bases

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