1. Quiz 1. Write the molecular formula of the following compounds: 3. Write the meaning of the following prefixes: Mono-, Tetra-, Penta-, Hexa-, Octa-, Deca- Sodium Chloride Calcium Phosphate Copper (II) Chloride Nitric Oxide Hydrocyanic acid Potassium Permagnate Sodium Borate Decahydrate CuSO 4 ． 5H 2 O 2. Write the name of the following compounds: Na 2 ClO 2 CuSO 3 0. How many ways are commonly used for denoting a molecule/ion?

2. Answer 1. Write the molecular formula of the following compounds: 3. Write the meaning of the following prefixes: Mono-1, Tetra-4, Penta-5, Hexa-6, Octa-8, Deca-10 Sodium Chloride NaCl Calcium Phosphate Ca 3 (PO4) 2 Copper (II) Chloride CuCl 2 Nitric Oxide HNO 2 Hydrocyanic acid HCN Potassium Permagnate KMnO 4 Sodium Borate Decahydrate NaBO 3.10H 2 O CuSO 4 ． 5H 2 O 2. Write the name of the following compounds: Na 2 ClO 2 CuSO 3 0. How many ways are commonly used for denoting a molecule/ion? (1) Molecular structural formula, (2) Ball-and-stick model, (3) Tube structure, (4) Space-filling representation. Copper (II) sulfate Sodium chlorite Copper (II) sulfate pentahydrate

3. Chapter 3 Chemical Reactions The carbon cycle is evident in fossils like this one, which are found in limestone, a form of calcium carbonate. The carbon atoms in limestone were once part of carbon dioxide molecules in the atmosphere. They were then taken up in the shells of marine organisms. When the organisms died, the shells settled to the bottom of the ocean and became compacted into limestone. Millions of years later, we dig up the limestone and use it to construct buildings. Some of the limestone is also heated to make quicklime in a process that releases the carbon atoms once again to the atmosphere as carbon dioxide. The process that brings about a chemical change

4. Skeletal Equation Reactants  Products Na+H 2 O  NaOH+H 2 sodium+water  sodium hydroxide+hydrogen Skeletal equation Staring materials Substances formed in a chemical reaction A reagent is a reactant only when it is being used in a particular reaction.

5. Chemical Equations Law of Conservation of Mass: Atoms are neither created nor destroyed in a chemical reaction. Balanced expression of chemical reaction=chemical equation 2Na+2H 2 O  2NaOH+H 2 Na+H 2 O  NaOH+H 2 Stoichiometric coefficients which give the molar ratios of the reactants and products Molecules

6. Reaction Conditions States: gas(g), liquid(l), aqueous(aq), solid(s) 2Na(s)+2H 2 O(l)  2NaOH(aq)+H 2 (g) Temperature CaCO 3 (s) CaO(s)+CO 2 (g) High temperature Other conditions: pressure, reaction time, catalysts …

7. Balancing Chemical Equations H2+O2H2OH2+O2H2O H2+O22H2OH2+O22H2O 2H2+O22H2O2H2+O22H2O 2H 2 (g)+O 2 (g)  2H 2 O(l) H 2 +O 2  H 2 O 2 2H+O  H 2 O H 2 +1/2O 2  H 2 O Change the stoichiometric coefficients only! ! Danger!

8. Balancing A Chemical Reaction C 4 H 10 +O 2  CO 2 +H 2 O C 4 H 10 +O 2  4CO 2 +5H 2 O C 4 H 10 +(13/2)O 2  4CO 2 +5H 2 O 2C 4 H 10 +13O 2  8CO 2 +10H 2 O 2C 4 H 10 (g)+13O 2 (g)  8CO 2 (g)+10H 2 O(l) !

9. Figure 3.6 When solutions of silver nitrate and potassium chromate are mixed, a precipitate of red silver chromate, Ag 2 CrO 4, forms. Precipitation reaction potassium chromate+silver nitrate  Silver chromate+potassium nitrate K 2 CrO 4 (aq)+2AgNO 3 (aq)  Ag 2 CrO 4 (s)+2KNO 3 (aq) Insoluble substance Soluble substance

10. Figure 3.7 These two beakers contain solutions with different concentrations of the same solute. The lighter color of the solution on the left (a) shows that the solute is less concentrated than in the solution on the right (b). In the molecular-level view, we see that there are more solute particles in a given volume of the more concentrated solution. solvent Solution=solvent+solute Dissolve

11. Concentration The amount of solute molecules in a given volume of solution

12. Figure 3.8 Sodium chloride consists of sodium ions and chloride ions. When it is added to water (left), the ions separate and spread throughout the solvent (right). The solution consists of water molecules, sodium ions, and chloride ions. There are no NaCl molecules present at any stage. The overlays show only the solute. Electrolyte: a substance that dissolves to give a solution that contains ions. Strong electrolytes: mostly ions. Weak electrolytes: mostly molecules Nonelectrolytes: no ions Hydration of ions

13. Nonelectrolyte

15. Figure 3.9 Pure water is a poor conductor of electricity, as shown by the almost imperceptible glow of the bulb in the circuit (a). However, when ions are present, as in an electrolyte solution, the solution does conduct. The ability of the solution to conduct is low if it is a weak electrolyte (b) but significant if it is a strong electrolyte (c), even if the solute concentration is the same in each case.

16. Figure 3.10 In a nonelectrolyte solution, the solute remains as molecules and does not break up into ions. Methanol, CH 3 OH, is a nonelectrolyte and is present as molecules when it is dissolved in water.

17. Figure 3.11 The formation of a silver chloride precipitate occurs immediately as silver nitrate solution is added to a solution of sodium chloride. AgNO 3 (aq)+NaCl(aq)  AgCl(s)+NaNO 3 (aq) Strong electrolyte

18. Figure 3.12 A series of scenes in a solution of sodium chloride. A sodium ion and a chloride ion move together, linger near each other for a time because of the attraction of their opposite charges, and then move apart. The loose, transient association of oppositely charged ions is called an ion pair. The solution is shown both with solvent molecules, for realism, and without, for clarity.

19. Figure 3.13 In water, ions are hydrated; that is, they are surrounded by a cluster of water molecules bound loosely to the ion. Note that a hydrated cation (a) is surrounded by water molecules oriented so that the O atom is closest to the ion, whereas a hydrated anion (b) has water molecules attached through their hydrogen atoms. The number of hydrating molecules depends on the size of the ion, but for most ions it is approximately six. Hydration of ions

20. Figure 3.14 In this precipitation reaction, yellow lead(II) chromate is formed when lead(II) nitrate solution is added to a solution of potassium chromate. Pb(NO 3 ) 2 (aq)+K 2 CrO 4 (aq)  PbCrO 4 (s)+2KNO 3 (aq)

21. Quiz Explain the following concepts: (1) Electrolyte (2) Hydration What is the real meaning of “ aq ” in a chemical equation?

22. Net Ionic Equations AgNO 3 (aq)+NaCl(aq)  AgCl(s)+NaNO 3 (aq) Complete ionic equation spectator ions Net ionic equation

23. Figure 3.15 Two depictions of a precipitation reaction that results when the ions in two electrolyte solutions are mixed (left beakers). The top right beakers show the fate of all four types of ions. By imagining the ionic reaction without the spectator ions (bottom right beakers), we can focus on the essential process described by the net ionic equation.

24. Figure 3.16 How to write a net ionic equation. Write the balanced overall equation (top), Then show all ionic solutes as separate ions in the complete ionic equation (second line), and delete the spectator ions. The result is the net ionic equation (bottom).

26. Figure 3.17 Another example of a precipitation reaction. This time, a solution of mercury(I) nitrate is being added to a solution of potassium iodide, and the insoluble product, mercury(I) iodide, is precipitated. Notice that a yellow color forms first. Mercury(I) iodide has two solid forms. The yellow form precipitates first but is soon converted to the more stable orange form.

27. Figure 3.18 The shape of this shell is a result of a precipitation reaction in which the shellfish secreted calcium ions at certain points on its surface. The calcium ions reacted with carbonate ions in the surrounding water. The colors of the shell are due to iron impurities that were captured in the solid as it formed.

28. The Reactions of Acids and Bases HCl (hydrochloric acid) CH 3 COOH (Acetic acid) NaOH (sodium hydroxide) NH 4 OH(ammonium hydroxide) Arrhenius acid Arrhenius base

29. HCl (hydrochloric acid) (Almost completely ionized in aqueous solution) Strong/Weak Acids CH 3 COOH (Acetic acid) (incomplete ionized in aqueous solution)

30. NaOH (sodium hydroxide) (Almost completely ionized in aqueous solution) Strong/Weak Bases NH 4 OH(ammonium hydroxide) (Incompletely ionized in aqueous solution)

31. The strong acids and bases in water HBr(aq), HCl(aq), HI(aq), HNO 3 (aq), HClO 4 (aq), HClO 3 (aq), H 2 SO 4 (aq) Group 1 hydroxides, Alkaline earth metal hydroxides

32. Neutralization Base+Acid  Salt +Water +(Others) HCl(aq)+NaOH(aq)  NaCl(aq)+H 2 O(l) 2HNO 3 (aq)+Mg(OH) 2  Mg(NO 3 ) 2 (aq)+2H 2 O(l) Neutralization=proton transfer

33. Gas-Formation Reactions 2NaCl(s)+H 2 SO 4 (l)  Na 2 SO 4 (s)+2HCl(g) FeS(s)+2HCl(aq)  FeCl 2 (aq)+H 2 S(g) CaCO 3 (s)+H 2 SO 4 (aq)  CaSO 4 (s)+H 2 CO 3 ( aq)  H 2 O+CO 2 (g) !!! The reaction of acids with salts is a proton transfer reaction that may produce gas or a compound that decomposes into a gas.

34. Redox Reactions 6CO 2 (g)+6H 2 O(l)  C 6 H 12 O 6 (s)+6O 2 (g) (photosynthesis reaction) CH 4 (g)+2O 2 (g)  CO 2 (g)+2H 2 O(l) (Natural gas reaction) 2Mg(s)+O 2 (g)  2MgO(s) Mg(s)+Cl 2 (g)  MgCl(s) Zn(s)+2HCl(aq)  ZnCl 2 (aq)+H 2 (g) Anything in common?

35. Figure 3.26 An example of an oxidation reaction: magnesium burning brightly in air. Magnesium also burns brightly in water and in carbon dioxide; consequently, magnesium fires are very difficult to extinguish. 2Mg(s)+O 2 (g)  2MgO(s) Oxidized (reducing agent) Reduced (Oxidizing agent)

36. Figure 3.27 When bromine is poured on red phosphorus, a vigorous reaction takes place. In the reaction phosphorus is oxidized and bromine is reduced. P(s)+5Br(s)  PBr 5 (s) Oxidized (reducing agent) Reduced (Oxidizing agent)

37. They Are All Redox Reactions 6CO 2 (g)+6H 2 O(l)  C 6 H 12 O 6 (s)+6O 2 (g) (photosynthesis reaction) CH 4 (g)+2O 2 (g)  CO 2 (g)+2H 2 O(l) (Natural gas reaction) 2Mg(s)+O 2 (g)  2MgO(s) Mg(s)+Cl 2 (g)  MgCl(s) Zn(s)+2HCl(aq)  ZnCl 2 (aq)+H 2 (g)

38. Figure 3.28 The common oxidation numbers of main-group elements. Notice the tendency of elements in the same group to assume the same oxidation number. How many electrons you want?

39. Figure 3.29 : How to determine an oxidation number. Each atom is imagined to be a separate ion. Certain ions are assigned charges by using the rules in Toolbox 3.3, and the charge on the central atom is determined by considering the overall charge on the species. (a) Oxide “ions” in an oxoanion are given the charge of  2; because there are four oxygen atoms and the overall charge on the anion is  2, the charge on the central atom must be  6. (b) This molecule has three chlorine atoms with oxidation numbers of  1, an oxygen atom (  2), and a hydrogen atom (  1). The sum of these oxidation numbers is  4 and the overall charge on the molecule is 0. Thus, the central atom must have an oxidation number of  4.

40. Determine Oxidation Number SO 2 X+2(-2)=0  x=4 x+4(-2)=-2  x=6

41. Figure 3.30 When a strip of zinc is placed in a solution that contains Cu 2  ions, the blue solution slowly becomes colorless and copper metal is deposited on the zinc. In this redox reaction, the zinc metal is reducing the Cu 2  ions to copper and the Cu 2  ions are oxidizing the zinc metal to Zn 2  ions. (a) The reaction. (b) A visualization of the process.

42. Figure 3.31 (a) Copper reacts slowly with dilute nitric acid to give blue Cu 2  ions and the colorless gas nitric oxide, NO. (b) When copper reacts with concentrated nitric acid, nitrogen dioxide, NO 2, is produced instead of NO. The blue solution is turned green by this brown gas.

43. Figure 3.32 Aluminum reacts vigorously with hydrochloric acid to form soluble aluminum chloride and water.

44. Case Study 3 Astronauts on the space shuttle must change the canisters of lithium hydroxide daily. Here, Sidney Gutierrez carries out the task. Two canisters are used, and one is changed every 12 hours so that the capacity to remove carbon dioxide remains stable. CO 2 (g)+2LiOH  Li 2 CO 3 (s)+H 2 O(l) A Better Solution: 4KO 2 (s)+2CO 2 (g)  K 2 CO 3 (s)+3O 2 (g) CO 2 (g)+2H 2 (g)  C(s)+2H 2 O(l) 2H 2 O(l)  2H 2 (g)+O 2 (g) (Each element can be recovered and reused!)

45. Figure 3.33 The three main types of chemical reactions discussed in this chapter can be distinguished by the type of change taking place. (a) In a precipitation reaction, ions mix and one combination of ions is insoluble. (b) In a neutralization reaction, hydrogen ions are transferred from an acid to a base. (c) In a redox reaction, electrons are transferred from a reducing agent to an oxidizing agent.

46. Figure 3.34 : We can predict the products of a reaction by examining the reactants. (a) Two soluble salts may form a precipitate. (b) An acid and a hydroxide react to form a salt and water. (c) When two elements react, a redox reaction generally occurs. A metal and nonmetal react to form an ionic compound and two nonmetals react to form a molecular compound. (d) In combustion reactions, organic compounds react with oxygen to form carbon dioxide and water.

47. Three Most Important Types of Reactions Precipitation (Soluble salts exchange ions  ionic solids Proton transfer (Neutralization, Gas Formation) Electron transfer (Redox Reaction)

48. Assignment for Chapter 3 17,25,33,37,43,51,62