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Chemical Reactions

Chemical Reactions

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Chemical Reactions

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  1. Chemical Reactions Chapter 4

  2. The Periodic Table • Review • Arrangement is based on increasing ____ ____. • Groups or columns • Elements within a group have similar ____ and _____ properties. • Alkali metals (Group __), alkaline earth metals (Group __), halogens (Group __), and noble gases (Group ___). • s, p, d, and f blocks

  3. Properties of Metals, Nonmetals and Metalloids

  4. Properties of Metals, Nonmetals and Metalloids • The stair sep line separates the metals from the nonmetals • Metals are to the left of this line • ~80% of the elements are metals • Nonmetals are to the right of this line • Metalloids are on the line (in green) • Metalloids or semimetals exhibit properties intermediate between metals and metalloids. • For semiconductors such as silicon, the conductance increases with temperature Typical properties exhibited by metals and nonmetals are given in Tables 4-3 and 4-4.

  5. Properties of Metals, Nonmetals and Metalloids • Metallic character increases from top to bottom and decreases from left to right. • Nonmetallic character decreases from top to bottom and increases from left to right. • Best nonmetals are far to the right Less Metallic More Metallic Periodic Chart

  6. Aqueous Solutions • An aqueous solution exists when a solute is dissolved in _____. • Formula is followed by (__) • An electrolyte is a substance whose aqueous solution _____ ______ due to the movement of ions or charged particles. • An electrolyte produces ____ when dissolved in water. For example, NaCl dissolved in water to produce Na+ and Cl- ions.

  7. Aqueous Solutions • The more ions produced by the substance, the stronger the electrolyte • Strong electrolyte – substances that conduct electricity well in dilute aqueous solutions • Weak electrolyte – substances that conduct electricity poorly in dilute aqueous solution • Nonelectrolyte – substances that do not conduct electricity in aqueous solutions DEMO: NaCl, acetic acid, and sugar water

  8. Generation of Ions in Solution • Ions are generally produced from a substance in water by either dissociation or ionization. • Dissociation – an _____ compound separates into ions in solution • NaCl(s) • Ionization – a ______ compound separates or reacts with water to form ions in solution • HNO3(aq) + H2O(l)  H3O+(aq) + NO3-(aq)

  9. Generation of Ions in Solution • Strong electrolytes. These are soluble in water. • Strong acids • Strong bases • Most soluble salts • Weak electrolytes. • Weak acids • Weak bases

  10. Acids and Bases • An acid is a substance that____ a proton. • In aqueous solution, this proton combines with water to form H3O+. All acids are ________ substances. • Strong acids • Ionize almost completely to form ions in dilute aqueous solutions. • Original acid molecules largely do not exist in solution. • Ionization is near 100% • A list of strong acids is given in Table 4-5 (know these).

  11. Acids and Bases • Representing reactions between strong acids and water. • The double arrow indicates that the reaction proceeds in both directions. The reaction is ______. In reactions with double arrows, the limiting reactant is not all used up. • The longer arrow pointing to the right indicates the the reaction is product-favored. • Ionization is near ________, only a _____ amount of limiting reactant remains. • All strong acids react similarly in water. They are product-favored.

  12. Acids and Bases • Weak acids ionize only _____ in dilute aqueous solutions. • Ionization is generally less than 5%. • A list of common weak acids is in Table 4.6. • The reaction is also reversible, but the longer arrow points to the left indication that the reaction is ________. • Most of the limiting reactant is not used.

  13. Reversible Reactions • Acetic acid/H2O reaction • HCl/H2O reaction • Al(NO3)3/NaOH reaction

  14. Acids and Bases • Most acids are weak. • ______ acids are almost always weak. • The O-H bond is broken. • Demonstration with acetic acid, CH3COOH (show model). • Write the reaction for the following acids with water. • HI, H2SO4, and (COOH)2

  15. Acids and Bases • A base is a substance the produces____ ions in aqueous solutions. The ___ ions produced, the stronger the base. Know the strong bases in Table 4-7. • Generation of OH- ions in aqueous solutions • Dissociation of metallic hydroxides • These dissolve and dissociate in water. • What do the arrows indicate in the reaction?

  16. Acids and Bases • Generation of OH- ions in aqueous solutions • Bases that ionize in water NH3(aq) + H2O(l) NH4+(aq) + OH-(aq) • These substances ionize in order to produce OH- • Mostly molecular substances • Most molecular substances that ionize to produce OH- are weak bases. • What do the arrows indicate about the limiting reactant?

  17. Acids and Bases • Some ionic hydroxides are essentially insoluble in water. Cu(OH)2 CuOH+ + OH- • Only a small amount of OH- ions are generated due to solubility. DEMO: NaOH, NH3, and an insoluble hydroxide in water. Classify these bases.

  18. Acids and Bases – A Summary • Strong acids and bases • The substance largely dissolves or reacts to produce ions (H3O+ or OH-) • Very small amount of limiting reactant remaining • Large arrow points toward the product side. • Weak acids and bases • Very small amount of the substance dissolves or reacts with water to produce ions (H3O+ or OH-) • Large amount of limiting reactant remaining • Large arrow points toward the reactant side

  19. Determining if a Compound is Soluble • Compounds whose solubility in water is less than 0.02 mol/L are generally classified as _______. • DEMO: Mg(OH)2 and NaCl • Most substance are not infinitely soluble. • There is a limit to how much solid will dissolve in water.

  20. Determining if a Compound is Soluble

  21. Determining if a Compound is Soluble • Determine if the below compounds are soluble in water. • K2SO4, PbCl2, MgCO3, and NaOH • Acids and bases that ionize or dissociate only slightly can still be soluble in water. • CH3COOH and HF You do not have to memorize the solubility rules. Make sure, however, that you can use the table.

  22. Reaction in Aqueous Solutions • Formation of insoluble ionic compounds • Pb(NO3)2(aq) + NaI(aq)  ????? • Will a solid compound form? How can you determine this?? • Refer back to the solubility rules. Is there a combination that would be insoluble? • Al(NO3)3(aq) + NaOH(aq)  ????? • NaCl(aq) + Ni(NO3)2(aq)  ?????

  23. Representing the Reactions in Aqueous Solutions • Formula unit equations • All complete formulas are shown 2AgNO3(aq) + Cu(s)  2Ag(s) + Cu(NO3)2(aq) DEMO: Reaction Write the formula unit equations for the previous reactions • Total ionic equations • Formulas illustrate the predominant form that is present in the aqueous solutions. • All species are still present. 2Ag+(aq) + 2NO3-(aq) + Cu(s)  2Ag(s) + Cu2+(aq) + 2NO3-(aq) Write the total ionic equations for the previous reactions

  24. Representing the Reactions in Aqueous Solutions • Net ionic equations • Only the species that react (i.e. change) are shown. Eliminate ions that don’t react. These are termed as spectator ions. 2Ag+(aq) + Cu(s)  2Ag(s) + Cu2+(aq) Write the net ionic equations for the reactions done previously.

  25. Oxidation Numbers • Reactions that involve the formation of ionic compounds include the transfer of electrons • NaCl and CaBr2 • For these ionic compounds, oxidation numbers are used to keep track of electron transfers. • The oxidation number (or oxidation state) of the ions in these compounds is simply the charge on the respective ion. • NaCl and CaBr2

  26. Oxidation Numbers • In molecular compounds, oxidation numbers are also assigned in order to aid in writing formulas and balancing equations. • Oxidation numbers are assigned on a per atom basis • Treat the rules in order of decreasing importance Page 138 (rules 1-8)

  27. Useful Rules in Assigning Oxidation Numbers • The oxidation number of any free, uncombined element is zero • The oxidation number of an element in a simple (monatomic) ion is the charge on the ion • In the formula for any compound, the sum of the oxidation numbers of all elements in the compound is zero. In a polyatomic ion, the sum of the oxidation numbers of the constituent elements is equal to the charge on the ion Table 4-10

  28. Assigning Oxidation Numbers • NaNO3 • K2Sn(OH)6 • H3PO4 • SO32- • Cr2O72-

  29. Naming Binary Ionic CompoundsIUPAC • Generally a metal combines with a nonmetal • Less electronegative element first • Use element’s full name • More electronegative element second • Stem (drop last part) and add ‘ide’ • Name the following ionic compounds LiBr, MgCl2, and Al2O3 • Indicate the charge on metals that can have multiple charge states (transition metals and Groups IIIA (except Al), IVA, and VA) • SnO, SnO2, and FeBr3

  30. Naming Ionic Compounds with Polyatomic Ions • Naming is very similar to the binary ionic compounds • The polyatomic ion charge and name is taken from Table 2-3 and Table 4-11. • Name the following ionic compounds • Al2(SO4)3, Ca(NO3)2, and (NH4)2SO4

  31. Naming Binary Molecular Compounds • Nonmetal bonded to a nonmetal • The first element listed (most metallic) is named first • Receives full name. • The second element listed in the formula is named second • Drop the last portion of the name and add ‘ide’ • Relative amounts of each element is indicated by a prefix • 2 – di, 3 – tri, 4 – tetra, 5 – penta, 6 – hexa

  32. Naming Binary Molecular Compounds • Let’s name some binary molecular compounds • SO3, OF2, and P4O6 • The minimum number of prefixes are used to name the compound unambiguously • mono, as a prefix, is not used • Exception is CO, carbon monoxide • The final ‘a’ in a prefix is omitted when the element begins with ‘o’ • As4O6

  33. Naming Binary Acids • Compounds in which H is bonded to a Group VIA element other than O or to a Group VIIA element • The pure compounds are named as typical binary compounds • Aqueous solutions of the compounds are named by adding ‘hydro’ as a prefix and ‘ic’ as a suffix. • For the suffix, ‘ic’ the last portion of the element is dropped

  34. Naming Binary Acids

  35. Naming Ternary Acids • The acids are composed of H, O, and one more element (usually a nonmetal) • H2SO4 • The ternary acids differ in the ____ of oxygen atoms contained in the acid • H2SO4 and H2SO3 • The nonmetal must be able to have more than one oxidation state • Two of the acids are chosen as a basis (reference) • The acid with the higher oxidation state number on the nonmetal or higher number oxygen atoms has a suffix ‘ic’ • The acid with the lower oxidation state on the nonmetal or smaller number of oxygen atoms has a suffix ‘ous’

  36. Naming Ternary Acids

  37. Naming Ternary Acids • Ternary acids that have one less O atom than the ‘ous’ acid is formed by adding the prefix ‘hypo’ and the suffix ‘ous’ • H3PO2 – hyposphosporous acid • What is the charge on the phosporus? • Name HClO and determine the charge on Cl • HClO3 is chloric acid and HClO2 is chlorous acid

  38. Naming Ternary Acids • Acids containing one more oxygen atom than the ‘ic’ acid are named by adding ‘per’ as a prefix and ‘ic’ as a suffix • HClO4 – perchloric acid • Name HIO4 and determine the oxidation state on I • HIO3 is iodic acid

  39. Naming Ternary Salts • These compounds are formed by replacing the hydrogen in a ternary acid with another ion that is usually a metal cation (occasionally NH4+) • If the ternary acid ended in ‘ic’, the ‘ic’ is replaced with ‘ate’ • If the ternary acid ended in ‘ous’, the ‘ous’ is replaced with ‘ite’ • The ‘per’ and ‘hypo’ prefixes are retained with the salts

  40. Naming Ternary Salts

  41. Naming Ternary Salts • Hydrogen is included in naming the salts and the ions • Name the following ternary salts and ions • NaHSO4 and KH2PO4 • H2PO2-1 and SO4-2 • Page 145

  42. Naming Ternary Acids and Salts • Page 145, Problem-Solving Tip Problems

  43. Redox Reactions • If there is a change in the oxidation number of an element in the reaction, the reaction can be classified as redox. • 4Al(s) + 3O2(g)  2Al2O3(s) Assign oxidation states to all atoms. Which element gained a greater oxidation state and which element gained a lower oxidation state?

  44. Redox Reactions • Oxidation – an ______ in oxidation state or apparent ____ of electrons. • Reduction – a ______ in oxidation state or apparent ____ of electrons. Oxidation and reduction occur _________ in a reaction (hence redox) • Oxidizing agents – species that oxidize other substances • Gain electrons and are reduced • Reducing agents – species that reduce other substances • Lose electron and are oxidized

  45. Redox Reactions • In the first reaction with Al and O2, identify the substances that are reduced and oxidized. Identify the oxidizing agent and reducing agent. • 3Zn(s) + 2CoCl3(aq)  3ZnCl2(aq) + 2Co(s) Is this a redox reaction? Write the net ionic equaiton. Identify oxidation states, the reduced and oxidized species, and the reducing and oxidizing agents.

  46. Redox Reactions • Al(NO3)3(aq) + 3NaOH(aq)  Al(OH)3(s) + 3NaNO3(aq) • 4KClO3(s)  KCl(s) + KClO4(s) In the last reaction, the same element (Cl) is oxidized and reduced. This is called a ______ reaction. Obtain the information outlined in the previous slide for each reaction.

  47. Combination Reaction • In this type reaction, two simpler substances combine to form a compound of the combined simpler substances. • 2Mg(s) + O2(g)  2MgO(s) • P4(s) + 10Cl2(g)  4PCl5(s) • CaO(s) + CO2(g)  CaCO3(s) Which of these reactions is a redox? When combining _______ the reaction is always redox.

  48. Decomposition Reactions • In this type of reaction, a compound decomposed in simpler substances. • 2H2O(l)  2H2(g) + O2(g) • 2KClO3(s)  2KCl(s) + 3O2(g) • (NH4)2Cr2O7(s)  Cr2O3(s) + 4H2O(g) + N2(g) Determine which reactions are redox? If the reaction involves the formation of a ______, the reaction is always redox.

  49. Displacement Reactions • In this reaction, one element displaces another in a compound. These reactions are always _____. 2AgNO3(aq) + Cu(s)  2Ag(s) + Cu(NO3)2(aq) • Do you recall the total ionic and net ionic equations? • This is a displacement reaction where a more active metal displaces a less active metal in aqueous solution. • More active metal + salt of less active metal  Less active metal + salt of more active metal