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Chemical Reactions

Chemical Reactions. Unit 4. 4.1 Introduction for Reactions. A physical change occurs when a substance undergoes a change in properties, but not a change in composition

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Chemical Reactions

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  1. Chemical Reactions Unit 4

  2. 4.1 Introduction for Reactions • A physical change occurs when a substance undergoes a change in properties, but not a change in composition • Changes in the phase of a substance (solid, liquid, gas)or formation/separation of mixtures of substances are common physical changes

  3. Introduction for Reactions • A chemical change occurs when substances are transformed into new substances, typically with different compositions. • Possible evidence that a chemical change has occurred: • Production of heat or light • Formation of a gas • Formation of a precipitate • Color change

  4. 4.2 Net Ionic Equations • All physical and chemical processes can be represented symbolically by balanced equations • Chemical equation represent chemical changes • These changes are the result of a rearrangement of atoms into new combinations

  5. Net Ionic Equations • Any representation of a chemical change must contain equal numbers of atoms of every element before and after the change occurred, demonstrating that mass is conserved in chemical reactions • Balanced molecular, complete ionic, and net ionic equations are differing symbolic forms used to represent a chemical reaction • The form used depends on the context

  6. Describing Reactions in Solution • Formula equation: gives the overall balanced reaction stoichiometry (like in first-year chemistry) • Complete ionic equation: represents all reactants and products as ions (if they all are strong electrolytes) • Net ionic equation: includes only those solution components undergoing a change, does not include spectator ions (do not include any ions that occur on both sides)

  7. Spectator Ions • Spectator ions: ions present in solution that do not participate directly in a reaction • They begin as aqueous ions and end as aqueous ions

  8. Example • Write the formula equation, the complete ionic equation, and the net ionic equation for: • Aqueous potassium chloride is added to aqueous silver nitrate to form a silver chloride precipitate and aqueous potassium nitrate. • Formula equation: KCl(aq) + AgNO3(aq)AgCl(s) + KNO3(aq) • Complete ionic equation: K+(aq) + Cl-(aq) + Ag+(aq) + NO3-(aq)AgCl(s) + K+(aq) + NO3-(aq) • Net ionic equation: Cl-(aq) + Ag+(aq)AgCl(s)

  9. Practice • Write the formula equation, the complete ionic equation, and the net ionic equation for: • Aqueous potassium hydroxide is mixed with aqueous iron(III) nitrate to form a precipitate of iron(III) hydroxide and aqueous potassium nitrate

  10. Answers to the Previous Slide • Formula equation: 3KOH(aq) + Fe(NO3)3(aq)Fe(OH)3(s) + 3KNO3(aq) • Complete ionic equation: 3K+(aq) + 3OH-(aq) + Fe3+(aq) + 3NO3-(aq)Fe(OH)3(s) + 3K+(aq) + 3NO3-(aq) • Net ionic equation: 3OH-(aq) + Fe3+(aq)Fe(OH)3(s)

  11. 4.3 Representations of Reactions • Balanced chemical equations in their various forms can be translated into symbolic particulate representations

  12. 4.4 Physical and Chemical Changes • Processes that involve the breaking and/or formation of chemical bonds are typically classified as chemical processes • Processes that involve only changes in intermolecular interactions, such as phase changes, are typically classified as physical processed

  13. Physical and Chemical Changes • Sometimes physical processes involve the breaking of chemical bonds • The dissolution of a salt in water could be plausibly argued as either a physical or chemical process • The dissolution involves breaking of ionic bonds and the formation of ion-dipole interactions between ions and solvent

  14. 4.5 Stoichiometry • Because atoms must be conserved during a chemical process, it is possible to calculate product amounts by using known reactant amounts, or to calculate reactant amounts given known product amounts • Coefficients of balanced chemical equations contain information regarding the proportionality of the amounts of substances involved in the reaction

  15. Stoichiometry • The values can be used in chemical calculations involving the mole concept • Stoichiometric calculations can also be combined with the ideal gas law and calculations involving molarity to quantitatively study gases and solutions

  16. Stoichiometry of Precipitation Reactions • Problem solving strategy: • Identify the species present and determine what reaction occurs • Write the balanced net ionic equation • Calculate the moles of reactants • Determine which reactant is limiting • Calculate the moles of product(s) as required • Convert to grams or other units, as required

  17. Example • Calculate the mass of solid NaCl that must be added to 1.50 L of a 0.100 M AgNO3 solution to precipitate all the Ag+ ions in the form of AgCl. • Ions present? Ag+ NO3- Na+ Cl- • Net ionic equation: Ag+(aq) + Cl-(aq)AgCl(s) • Moles of Ag+ present in the solution? (1.50 L) (0.100 mol Ag+) = 0.150 mol Ag+ • mol of Cl- to react with all of the mol Ag+? 1:1 mol ratio so 0.150 mol of Cl- and 0.150 mol NaCl • 0.150 mol NaCl = 8.77 g NaCl

  18. Gas Stoichiometry • Solve for the volume of one mol of gas at STP: • V = nRT V = (1) (0.08206) (273) = P 1 • 22. 4 L • This is the molar volume of a gas at STP • The ideal gas law can be used to determine moles in a more complicated stoichiometry problem

  19. Solution Stoichiometry • You can calculate the number of moles from molarity and volume (in liters) • n = ML • You may have to consider limiting and excess reactants

  20. 4.6 Introduction to Titration • Titrations may be used to determine the concentration of an analyte in solution • The titrant has a known concentration of a species that reacts specifically and quantitatively with the analyte • You must have a balanced equation to calculate the concentration of the analyte, and you may have to review solution stoichiometry

  21. Introduction to Titration • The equivalence point of the titration occurs when the analyte is totally consumed by the reacting species in the titrant • The equivalence point is often indicated by a change in a property (such as color) that occurs when the equivalence point is reached • This observable event is called the endpoint of the titration

  22. Reaction of HCl with NaOH

  23. 4.7 Types of Chemical Reactions • Acid-base reactions involve transfer of one or more protons (H+) between chemical species • Oxidation-reduction (redox) reactions involve transfer of one or more electrons between chemical species, as indicated by changes in oxidation numbers of the involved species • In a redox reaction, electrons are transferred from the species that is oxidized to the species that is reduced

  24. Assigning Oxidation States • An atom in an element is zero, including allotropes (N2, S8, Mg) • A monatomic ion is the same as its charge • Fluorine is always -1 (unless it is an atom) • Oxygen is usually -2, except in peroxides (-1) or when bonded with fluorine (+2) • Hydrogen is +1 when bonded with nonmetals, and -1 when bonded with metals

  25. More Rules • The more-electronegative element in a binary molecular compound is assigned the number equal to the negative charge it would have as an anion, the other gets the positive charge it would get as a cation (this is what you do when you are stumped) • The sum of the oxidation states must be zero for an electrically neutral compound • For a polyatomic ion, the sum of the oxidation states must equal the charge of the ion

  26. Examples • Assign oxidation states to all atoms: • CO • Start with oxygen because it is usually -2 • To achieve zero, carbon must be +2 • This does not mean that carbon is always +2 • CO2 • Start with oxygen as -2, but you have two of them to consider • carbon is +4 while oxygen is -2

  27. SF6 • fluorine is always -1, so sulfur is +6 • NO3- • oxygen is -2, since the sum must equal -1, nitrogen is +5 • H2 • hydrogen is 0 • C2H5OH (ethyl alcohol, AKA ethanol) • oxygen is -2, hydrogen is +1, so carbon is -2 • Cl- chloride is -1

  28. Types of Chemical Reactions • Combustion is an important subclass of oxidation-reduction reactions, in which a species reacts with oxygen gas • In the case of hydrocarbons, carbon dioxide and water are products of complete combustion • Hydrocarbon: any of a class of organic chemical compounds containing only carbon and hydrogen

  29. Types of Chemical Reactions • Precipitation reactions frequently involve mixing ions in aqueous solution to produce an insoluble or sparingly soluble ionic compound • All sodium (Na+), potassium (K+), ammonium (NH4+), and nitrate (NO3-) salts are soluble in water • Gravimetric analysis is used to analyze the precipitate formed • The precipitate is frequently separated from the solution using filtration

  30. 4.8 Introduction to Acid-Base Reactions • By definition, a Brønsted-Lowry acid is a proton donor and a Brønsted-Lowry base is a proton acceptor • Only in aqueous solutions, water plays an important role in many acid-base reactions, as its molecular structure allows it to accept protons from and donate protons to dissolved species

  31. Introduction to Acid-Base Reactions • When an acid or base ionizes in water, the conjugate acid-base pairs can be identified and their relative strengths compared

  32. Conjugate Base • Conjugate base: the species that remains after a Brønsted-Lowry acid has given up a proton • Acid conjugate base

  33. Conjugate Acid • Conjugate acid: the species that is formed when a Brønsted-Lowry base gains a proton • base conjugate acid

  34. Acid-Base Reactions • Usually acid-base reactions are equilibrium systems • The forward and reverse reactions occur • They involve conjugate acid-base pairs CH3COOH(aq) + H2O(l) H3O+(aq) + CH3COO -(aq) acid base conjugate conjugate acid base

  35. Linking the Pairs

  36. Other Representations

  37. Strength of Conjugate Acids and Bases • The stronger an acid is, the weaker its conjugate base • The stronger a base is, the weaker its conjugate acid

  38. 4.9 Oxidation-Reduction (Redox) Reactions: Balancing • Balance the oxidation half-reaction and the reduction half-reaction • Balance the number of electrons lost and gained • Add the balanced half-reactions and cancel species that appear on both sides of the arrow • Check to see that both the number of atoms and the charges are balanced

  39. Example Problem • Balance the following and check the answer: • Fe2+ + Cl2 Fe3+ + Cl- • Oxidation half-reaction: Fe2+ Fe3+ + e- • The oxidation half-reaction is balanced • Reduction half-reaction: Cl2+ 2e- Cl- • Balanced reduction : Cl2 + 2e- 2Cl- • Balance electrons: 2Fe2+ 2Fe3+ + 2e- • Add: 2Fe2+ + Cl2 + 2e-2Fe3+ + 2e-+ 2Cl- • Cancel species: 2Fe2+ + Cl22Fe3+ + 2Cl-

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