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0. How many ways are commonly used for denoting a molecule/ion?. Quiz. Sodium Chloride Calcium Phosphate Copper (II) Chloride Nitric Oxide Hydrocyanic acid Potassium Permagnate Sodium Borate Decahydrate. 1. Write the molecular formula of the following compounds:.

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  1. 0. How many ways are commonly used for denoting a molecule/ion? Quiz Sodium Chloride Calcium Phosphate Copper (II) Chloride Nitric Oxide Hydrocyanic acid Potassium Permagnate Sodium Borate Decahydrate 1. Write the molecular formula of the following compounds: 2. Write the name of the following compounds: CuSO4．5H2O Na2ClO2 CuSO3 3. Write the meaning of the following prefixes: Mono-, Tetra-, Penta-, Hexa-, Octa-, Deca-

  2. 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. Answer Sodium Chloride NaCl Calcium Phosphate Ca3(PO4)2 Copper (II) Chloride CuCl2 Nitric Oxide HNO2 Hydrocyanic acid HCN Potassium Permagnate KMnO4 Sodium Borate Decahydrate NaBO3.10H2O 1. Write the molecular formula of the following compounds: 2. Write the name of the following compounds: CuSO3 CuSO4．5H2O Na2ClO2 Copper (II) sulfate Sodium chlorite Copper (II) sulfate pentahydrate 3. Write the meaning of the following prefixes: Mono-1, Tetra-4, Penta-5, Hexa-6, Octa-8, Deca-10

  3. Chapter 3 Chemical Reactions The process that brings about a chemical change 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.

  4. Skeletal Equation ReactantsProducts Skeletal equation Substances formed in a chemical reaction Staring materials A reagent is a reactant only when it is being used in a particular reaction. sodium+watersodium hydroxide+hydrogen Na+H2ONaOH+H2

  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 Na+H2ONaOH+H2 2Na+2H2O2NaOH+H2 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)+2H2O(l)2NaOH(aq)+H2(g) Temperature CaCO3(s)CaO(s)+CO2(g) High temperature Other conditions: pressure, reaction time, catalysts …

  7. Balancing Chemical Equations H2+O2H2O ! Danger! H2+O22H2O Change the stoichiometric coefficients only! 2H2+O22H2O 2H2(g)+O2(g)2H2O(l) H2+O2H2O2 2H+OH2O H2+1/2O2H2O

  8. Balancing A Chemical Reaction C4H10+O2CO2+H2O C4H10+O24CO2+5H2O C4H10+(13/2)O24CO2+5H2O 2C4H10+13O28CO2+10H2O 2C4H10(g)+13O2(g)8CO2(g)+10H2O(l) !

  9. Figure 3.6 When solutions of silver nitrate and potassium chromate are mixed, a precipitate of red silver chromate, Ag2CrO4, forms. Precipitation reaction Insoluble substance potassium chromate+silver nitrate Silver chromate+potassium nitrate K2CrO4(aq)+2AgNO3(aq)Ag2CrO4(s)+2KNO3(aq) 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 Dissolve Solution=solvent+solute

  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. Hydration of ions Electrolyte: a substance that dissolves to give a solution that contains ions. Strong electrolytes: mostly ions. Weak electrolytes: mostly molecules Nonelectrolytes: no ions

  13. Nonelectrolyte

  14. 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, CH3OH, 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. Strong electrolyte AgNO3(aq)+NaCl(aq)AgCl(s)+NaNO3(aq)

  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(NO3)2(aq)+K2CrO4(aq)PbCrO4(s)+2KNO3(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 AgNO3(aq)+NaCl(aq)AgCl(s)+NaNO3(aq) Complete ionic equation Net ionic equation spectator ions

  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).

  25. 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.

  26. 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.

  27. The Reactions of Acids and Bases • HCl (hydrochloric acid) • CH3COOH (Acetic acid) • NaOH (sodium hydroxide) • NH4OH(ammonium hydroxide) Arrhenius acid Arrhenius base

  28. Strong/Weak Acids HCl (hydrochloric acid)(Almost completely ionized in aqueous solution) CH3COOH (Acetic acid) (incomplete ionized in aqueous solution)

  29. Strong/Weak Bases • NaOH (sodium hydroxide) (Almost completely ionized in aqueous solution) • NH4OH(ammonium hydroxide)(Incompletely ionized in aqueous solution)

  30. The strong acids and bases in water • HBr(aq), HCl(aq), HI(aq), HNO3(aq), HClO4(aq), HClO3(aq), H2SO4(aq) • Group 1 hydroxides, Alkaline earth metal hydroxides

  31. Neutralization • Base+Acid Salt +Water +(Others) • HCl(aq)+NaOH(aq)NaCl(aq)+H2O(l) • 2HNO3(aq)+Mg(OH)2Mg(NO3)2(aq)+2H2O(l) Neutralization=proton transfer

  32. Gas-Formation Reactions !!! • 2NaCl(s)+H2SO4(l)Na2SO4(s)+2HCl(g) • FeS(s)+2HCl(aq)FeCl2(aq)+H2S(g) • CaCO3(s)+H2SO4(aq)CaSO4(s)+H2CO3(aq)H2O+CO2(g) The reaction of acids with salts is a proton transfer reaction that may produce gas or a compound that decomposes into a gas.

  33. Redox Reactions 6CO2(g)+6H2O(l)C6H12O6(s)+6O2(g) (photosynthesis reaction) CH4(g)+2O2(g)CO2(g)+2H2O(l) (Natural gas reaction) 2Mg(s)+O2(g)2MgO(s) Mg(s)+Cl2(g)MgCl(s)Zn(s)+2HCl(aq)ZnCl2(aq)+H2(g) Anything in common?

  34. 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)+O2(g)2MgO(s) Oxidized (reducing agent) Reduced (Oxidizing agent)

  35. Figure 3.27When 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)PBr5(s) Oxidized (reducing agent) Reduced (Oxidizing agent)

  36. They Are All Redox Reactions 6CO2(g)+6H2O(l)C6H12O6(s)+6O2(g) (photosynthesis reaction) CH4(g)+2O2(g)CO2(g)+2H2O(l) (Natural gas reaction) 2Mg(s)+O2(g)2MgO(s) Mg(s)+Cl2(g)MgCl(s)Zn(s)+2HCl(aq)ZnCl2(aq)+H2(g)

  37. 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?

  38. 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.

  39. Determine Oxidation Number SO2 X+2(-2)=0x=4 x+4(-2)=-2x=6

  40. Figure 3.30 When a strip of zinc is placed in a solution that contains Cu2+ 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 Cu2+ ions to copper and the Cu2+ ions are oxidizing the zinc metal to Zn2+ ions. (a) The reaction. (b) A visualization of the process.

  41. Figure 3.31 (a) Copper reacts slowly with dilute nitric acid to give blue Cu2+ ions and the colorless gas nitric oxide, NO. (b) When copper reacts with concentrated nitric acid, nitrogen dioxide, NO2, is produced instead of NO. The blue solution is turned green by this brown gas.

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

  43. Case Study 3Astronauts 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. A Better Solution: 4KO2(s)+2CO2(g)K2CO3(s)+3O2(g) CO2(g)+2H2(g)C(s)+2H2O(l) 2H2O(l)2H2(g)+O2(g) (Each element can be recovered and reused!) CO2(g)+2LiOHLi2CO3(s)+H2O(l)

  44. 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.

  45. 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.

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

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

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