1. ACIDS and BASES  Acid – Base theories  Naming acids and bases  Oxides  Reactions and properties of acids and bases  Strengths of acids and bases

2. Acid and Base Theories 1) Arrhenius Theory An acid is a substance that gives H + ion, when dissolved in water. For example, hydrochloric acid reacts with water to form hydrogen ions which are transferred to a water molecule to form a hydronium ion (H 3 O + ). But simply the reaction is: HCl H + + Cl -

3. Acids which have one ionizable hydrogen atom per molecule are called monoprotic acids. Example: HNO 3 H + + NO 3 - Acids which have two ionizable hydrogen atom per molecule are called diprotic acids. Example: H 2 SO 4 H + + HSO 4 − HSO 4 − ⇌ H + + SO 4 2− Acids which have three ionizable hydrogen atom per molecule are called triprotic acids. Example: H 3 PO 4 ⇌ H + + H 2 PO 4 – H 2 PO 4 – ⇌ H + + HPO 4 2– HPO 4 2– ⇌ H + + PO 4 3–

4. A base is a substance that gives OH - ion, when dissolved in water. NaOH → Na + + OH − Ca(OH) 2 → Ca 2+ + 2OH - Reaction of NH 3 produce OH - : NH 3 + H 2 O → NH 4 + + OH - so it is a base.

5. Limitations of the Arrhenius theory Hydrochloric acid is neutralized by both sodium hydroxide solution and ammonia solution. In both cases, you get a colourless solution which you can crystallize to get a white salt - either sodium chloride or ammonium chloride. These are clearly very similar reactions. The full equations are: NaOH(aq) + HCl(aq) NaCl(aq) + H 2 O(l) NH 3 (aq) + HCl(aq) NH 4 Cl(s) In the sodium hydroxide case, hydrogen ions from the acid are reacting with hydroxide ions from the sodium hydroxide - in line with the Arrhenius theory. However, in the ammonia case, there don't appear to be any hydroxide ions!

6. You can get around this by saying that the ammonia reacts with the water, it is dissolved in to produce ammonium ions and hydroxide ions: NH 3 (aq) + H 2 O(l) NH 4 + (aq) + OH - (aq) This is a reversible reaction, and in a typical dilute ammonia solution, about 99% of the ammonia remains as ammonia molecules. Nevertheless, there are hydroxide ions there, and we can squeeze this into the Arrhenius theory. However, this same reaction also happens between ammonia gas and hydrogen chloride gas. NH 3 (g) + HCl(g) NH 4 Cl(s) In this case, there aren't any hydrogen ions or hydroxide ions in solution - because there isn't any solution. The Arrhenius theory wouldn't count this as an acid-base reaction, despite the fact that it is producing the same product as when the two substances were in solution.

7. Naming Acids and Bases A. Naming Acids: The name of the acid is determined based on the name of the anion, specifically, based on the ending of the anion name. The three possibilities are listed here: Anion Name Acid Name -ideHydro-ic acid -ite-ous acid -ate-ic acid

8. FluorideF-F- ChlorideCl - BromideBr - IodideI-I- SulfideS 2- NitrideN 3- SulfiteSO 3 2- NitriteNO 2 - ChloriteClO 2 - HypochloriteOCl - PhosphatePO 4 3- Hydrogen phosphateHPO 4 2- Dihydrogen phosphateH 2 PO 4 - NitrateNO 3 - SulfateSO 4 2- Hydrogen sulfateHSO 4 - PerchlorateClO 4 - ChlorateClO 3 - CarbonateCO 3 2- Common Anions

9. B. Naming Bases Simply use the normal rules for naming compounds; ionic or covalent depending on the elements in the compound. Example: NaOH: Sodium hydroxide Ca(OH) 2 : Calcium hydroxide NH 3 : Ammonia

10. Example: a) Name the following acids and bases: NaOH: H 2 SO 3 : H 2 S : H 3 PO 4 : NH 3 : HCN: Ca(OH) 2 : Fe(OH) 3 : H 3 P: Sodium hydroxide Sulfurous acid Hydrosulfuric acid Phosphoric acid Ammonia Hydrocyanic acid Calcium hydroxide Iron (III) hydroxide Hydrophosphoric acid

11. b) Write the formulas of the following acids and bases: Hydrofluoric acid: Hydroselenic acid: Carbonic acid: Lithium hydroxide: Nitrous acid: Cobalt (II) hydroxide: Sulfuric acid: Beryllium hydroxide: Hydrobromic acid: HF H 2 Se H 2 CO 3 LiOH HNO 2 Co(OH) 2 H 2 SO 4 Be(OH) 2 HBr

12. Oxides Nonmetal Oxides Metal Oxides CO 2, SO 2, SO 3 etc. show acidic properties (acid anhydride) CO, NO, N 2 O are neutral (have 1 oxygen atom in the formula) Na 2 O, BaO etc. show basic properties (basic anhydrides) Amphoteric metals show both basic and acidic properties such as Al and Zn

13. Acidic Property of Nonmetal Oxides The oxides of nonmetals are usually acidic except NO, N 2 O and CO (They are neutral) CO 2 + H 2 O H 2 CO 3 SO 2 + H 2 O H 2 SO 3 SO 3 + H 2 O H 2 SO 4 N 2 O 5 + H 2 O 2HNO 3 Cl 2 O + H 2 O 2HOCl P 4 O 10 + H 2 O 4H 3 PO 4 Monoxides of halogens are acidic such as Cl 2 O, Br 2 O. Oxides of some metals at high oxidation states show acidic properties such as Mn 2 O 7, CrO 3. Acidic nonmetal oxides react with bases to form salts. SO 3 + 2KOH K 2 SO 4 + H 2 O

15. ACID ANHYDRIDES Carbon dioxide dissolved in water is in equilibrium with carbonic acid: equilibrium CO 2 + H 2 O ⇌ H 2 CO 3 The hydration equilibrium constant at 25°C is K h = 1.70×10 −3 : hence, the majority of the carbon dioxide is not converted into carbonic acid and stays as CO 2 molecules.hydrationequilibrium constant

16. Basic Properties of Metal Oxides Oxides of metals are usually basic. Na 2 O + H 2 O 2NaOH BaO + H 2 O Ba(OH) 2 Some metal oxides can not dissolve in water but they can dissolve in acidic solutions. MnO + 2HCl MnCl 2 + H 2 O CrO + 2HCl CrCl 2 + H 2 O Basic oxides react with acids to form salts. CaO + H 2 SO 4 CaSO 4 + H 2 O

19. ANHYDRIDES

20. Amphoteric Oxides Oxides amphoteric metals are also amphoteric. Al 2 O 3 + HCl AlCl 3 + H 2 O Al 2 O 3 + 2NaOH + 3H 2 O 2NaAl(OH) 4 (sodium tetrahydroxoaluminate)

21. Properties and Reactions of Acids and Bases A.Properties of Acids: Are corrosive They taste sour They form solutions w/ pH less than 7 at 25°C. They turn litmus dye from blue to red Their aqueous solutions conduct electricity (electrolyte) They react with active metals to form salt and H 2 gas. Mg + 2HCl MgCl 2 + H 2

22. The acids which do not contain oxygen in their structures can not react with semi noble metals Cu, Hg, Ag.The oxy acids react with these metals producing gases other than H 2. Cu + 2H 2 SO 4  CuSO 4 + SO 2 + 2H 2 O 3Ag + 4HNO 3  3AgNO 3 + NO + 2H 2 O They react with metal carbonates and hydrogen carbonates(bicarbonate ion) to give a salt, water and carbon dioxide, which appears as effervescence (bubbles). Na 2 CO 3 + 2HCl NaCl + H 2 O + CO 2 CH 3 COOH (aq)+NaHCO 3 (aq)  NaCH 3 COO(aq) +H 2 O (l) + CO 2 ethanoic acid metalh hydrogen salt water carbon carbonate dioxide

23. They react with bases to form salts and water. HCl + NaOH  NaCl + H 2 O (neutralization) H + (aq) + OH - (aq)  H 2 O(l) (net ionic equation)

24. B. Properties of Bases They have bitter taste Aqueous solutions of bases, known as alkali, have a slippery feel. They turn the litmus dye from red to blue They react with fats in the skin to form soaps Their aqueous solutions conduct electricity (electrolyte) The most common bases are the oxides, hydroxides and carbonates of metals, but a number of other compounds, such as ammonia and amines (CH 3 NH 2 ) also act as bases.

25. They only react with amphoteric metals: Zn, Al Zn + 2NaOH  Na 2 ZnO 2 + H 2 2Al + 6 NaOH  2Na 3 AlO 3 + 3H 2 If they are soluble in water they give a solution with pH>7 (at 25 o C). They react with acids to form a salt. CaO (s) + 2 HCl (aq)  CaCl 2 (aq) + H 2 O (l) base acid salt water

26. Amphoteric metals can react with both acids and bases, such as Al, Zn, Sn, Pb, Cr Al + 6HCl AlCl 3 + 3H 2 2Al + 6NaOH 2Na 3 AlO 3 + 3H 2 Oxides and hydroxides of amphoteric metals are also amphoteric. Al 2 O 3 + HCl AlCl 3 + H 2 O Al 2 O 3 + 2NaOH + 3H 2 O 2NaAl(OH) 4 ZnO + 2 HCl ZnCl 2 + H 2 O ZnO + 2NaOH + H 2 O Na 2 Zn(OH) 4

27. Neutralization

28. Examples of Acids & Bases Acids HCl H 2 SO 4 HNO 3 Juices, Soda Carbonic acid Acetic acid NaOH Ca(OH) 2 KOH Soap, Ammonia NH 3, Baking Soda Amines Bases

29. 2.Bronsted-Lowry Theory A Bronsted-Lowry (BL) acid is defined as any substance that can donate a hydrogen ion (proton) and a Bronsted-Lowry base is any substance that can accept a hydrogen ion (proton). Thus, according to the BL definition, acids and bases must come in conjugate pairs. For example, consider acetic acid dissolved in water: CH 3 COOH + H 2 O CH 3 COO - + H 3 O + Conjugate acid-base pairs: 1.CH 3 COOH and CH 3 COO - 2.H 2 O and H 3 O + Act as an acid Act as a base Act as an acid Act as a base

30. Label Bronsted-Lowry acids and bases in the following reactions and show the direction of proton transfer. 1.H 2 O + H 2 O OH - + H 3 O + AcidBase Acid H+H+ 2. NH 3 + H 2 O NH 4 + + OH - AcidBaseAcidBase H+H+ When a Bronsted-Lowry acid has given up its proton, it is capable of getting back that proton and acting as a base. Conjugate base is what is left after an acid gives up a proton. The stronger the acid, the weaker the conjugate base. The stronger the base, the weaker the conjugate acid.

31. The relationship between the Bronsted-Lowry theory and the Arrhenius theory The Bronsted-Lowry theory doesn't go against the Arrhenius theory in any way - it just adds to it. Hydroxide ions are still bases because they accept hydrogen ions from acids and form water. An acid produces hydrogen ions in solution because it reacts with the water molecules by giving a proton to them.

32. The hydrogen chloride / ammonia problem This is no longer a problem using the Bronsted-Lowry theory. Whether you are talking about the reaction in solution or in the gas state, ammonia is a base because it accepts a proton (a hydrogen ion). If it is in solution, the ammonia accepts a proton from a hydronium ion: NH 3 (aq) + H 2 O(l) NH 4 + (aq) + OH - (aq) If the reaction is happening in the gas state, the ammonia accepts a proton directly from the hydrogen chloride: NH 3 (g) + HCl(g) NH 4 Cl(s)

33. 3. Lewis Theory A Lewis acid is a chemical compound, A, that can accept a pair of electrons from a Lewis base, B, that acts as an electron- pair donor, forming an adduct, AB. A + :B → A—B A Lewis base is also a Brønsted-Lowry base.

35. The Bronsted-Lowry theory says that they are acting as bases because they are combining with hydrogen ions. The reason they are combining with hydrogen ions is that they have lone pairs of electrons - which is what the Lewis theory says. The two are entirely consistent. But what about other similar reactions of ammonia or water, for example?

36. Ammonia reacts with BF 3 by using its lone pair to form a co-ordinate bond with the empty orbital on the boron.

37. Co-ordinate (dative covalent) bonding A covalent bond is formed by two atoms sharing a pair of electrons. The atoms are held together because the electron pair is attracted by both of the nuclei. In the formation of a simple covalent bond, each atom supplies one electron to the bond. A co-ordinate bond (also called a dative covalent bond) is a covalent bond (a shared pair of electrons) in which both electrons come from the same atom.

38. The reaction between ammonia and hydrogen chloride

39. Representing co-ordinate bonds In simple diagrams, a co-ordinate bond is shown by an arrow. The arrow points from the atom donating the lone pair to the atom accepting it.

40. Dissolving hydrogen chloride in water to make hydrochloric acid

41. Lewis acids Lewis acids are electron pair acceptors. In the above example, the BF3 is acting as the Lewis acid by accepting the nitrogen's lone pair. On the Bronsted-Lowry theory, the BF3 has nothing about it. Then what makes HCl a Lewis acid? Chlorine is more electronegative than hydrogen, and that means that the hydrogen chloride will be a polar molecule. The electrons in the hydrogen-chlorine bond will be attracted towards the chlorine end, leaving the hydrogen slightly positive and the chlorine slightly negative.

42. The lone pair on the nitrogen of an ammonia molecule is attracted to the slightly positive hydrogen atom in the HCl. As it approaches it, the electrons in the hydrogen-chlorine bond are repelled still further towards the chlorine. Eventually, a co-ordinate bond is formed between the nitrogen and the hydrogen, and the chlorine breaks away as a chloride ion.

43. Relative Strengths of acids and Bases The strength of an acid depends on how easily the proton, H + is lost or removed from an acid Two factors determine the acidic strength: 1.The polarity of the bond with H atom: The more polarized the bond is, the more easily the proton is removed and greater the acidic strength. 2.The size of the atom X (in HX): The greater the atom X, the weaker is the bond and greater the acidic strength.

44. Periodic Trends for Binary Acids: Down a group: Sizes of the atoms increase. HF HCl Acidic strength increases HBr HI Across a period: Polarity of the bond increases. CH 4 NH 3 H 2 O HF Acidic strength inreases.

45. Oxyacids: HOF HOCl Acidic strength decreases. H-O HOBr bond ionizes HOI more easily when the oxygen atom is bonded to a more electronegative atom.

46. For a series of oxyacids: HClOHClO 2 HClO 3 HClO 4 Acidity increases As the number of oxygen atoms increases, The oxidation number of central atom (Cl) increases. This increases the ionization of O-H bond. Therefore, acidic strength increases.

47. Polyprotic Acids and Their Anions: H 3 PO 4 H 2 PO 4 - HPO 4 2- H 2 CO 3 HCO 3 - Acidity decreases H 2 SO 4 HSO 4 -

48. Organic Acids Organic acids have carboxyl group (COOH). They are weak acids. Example: HCOOH: Formic acid CH 3 COOH: Acetic acid

49. Basic strength As the volume of the metal increases, it becomes easier to ionize the OH - ion and the basic strength increases. LiOH NaOH KOH Basic strength increases