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Chapter 4 Types of Chemical Reactions and Solution Stoichiometry

Chapter 4 Types of Chemical Reactions and Solution Stoichiometry. 4.1 Water, the Liquid of Life 4.2 Aqueous Solutions: Strong and Weak Electrolytes 4.3 The Composition of Solutions 4.4 Types of Chemical Reactions 4.5 Precipitation Reactions 4.6 Describing Reactions in Solution

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Chapter 4 Types of Chemical Reactions and Solution Stoichiometry

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  1. Chapter 4Types of Chemical Reactions and Solution Stoichiometry 4.1 Water, the Liquid of Life 4.2 Aqueous Solutions: Strong and Weak Electrolytes 4.3 The Composition of Solutions 4.4 Types of Chemical Reactions 4.5 Precipitation Reactions 4.6 Describing Reactions in Solution 4.7 Selective Precipitation 4.8 Stoichiometry of Precipitation Reactions 4.9 Acid-Base Reactions 4.10 Oxidation-Reduction Reactions 4.11 Balancing Oxidation-Reduction Equations (skip) 4.12 Simple Oxidation-Reduction Titrations (skip)

  2. Water, the Liquid of Life (part 1) Water is polar. In a water molecule, each oxygen atom is partially negative. Each hydrogen atom is partially positive. This phenomenon is called charge separation. . Oxygen is an electron hog. It does not share electrons well with other atoms. Oxygen takes more than its share of electron density, that is why it has a partial negative charge. The hydrogen atoms are stripped of much of their electron density and so carry a partial positive charge. In liquid water and especially in solid water, the molecules interact strongly with each other, with preferred orientations.

  3. Water, the Liquid of Life (part 2) • Water has all sort of strange and unusual properties.1) Its density decreases when it freezes (ice floats).2) It has a high boiling point and high heat of vaporization.3) It has high surface tension.4) It dissolves many salts (like sodium chloride) and polar molecules (like ethanol).5) It does not dissolve non-polar substances (oil and water don’t mix).6) It has high heat capacity7) In the presence of amphipaths, it readily forms compartments (like in cells).8) Where there is liquid water, there is life (I bet).

  4. Water, the Liquid of Life (part 3) Water dissolves NaCl. NaCl dissolves in water. To dissolve is to mix at the level of individual ions and molecules. A solution is a homogeneous mixture of two or more substances. A solution may exist in any phase. Water is the solvent The solvent is the component of a solution that is present in the greatest amount. NaCl is the solute The solute is the substance that is dissolved in a solution. Pairs of liquids that mix in any proportion are miscible. Liquids that do not mix are immiscible.

  5. Ice

  6. Classification of Molecules 1) Non-polar: No charge separation, every atom is neutral. Examples: graphite = Cbig, hexane = C8H18, or olive oil. 2) Polar: With charge separation, some atoms carry partial charge. Examples: water = H2O, ammonia = NH3, fructose = C6H12O6.3) Ionic Compounds: With Formal (unit) charge.

  7. Ionic Compounds (Salts) Examples: Na+Cl-, Na+Br-, Mg2+Cl-2, NH4+NO3-, CH3COO-Na+ Salts contain charged (unit charge, not partial charge) atoms or groups of atoms. Cations: Na+, K+, Mg2+ Ca2+, NH4+ Anions: Cl-, F-, SO42-, PO43-, CH3COO-

  8. Properties of Ionic Compounds: All ionic compounds are solids in the range of temperature in which water is a liquid. They have rigid lattices in which strong electrostatic forces (ionic bonds) hold the constituent ions in place. The high melting points of ionic compounds indicate that a good deal of energy must be supplied to destroy the lattice and produce a liquid (molten) form in which the ions move more freely. Solubility in Water: Water molecules are polar. When water interacts with an ionic compound such as NH4NO3, or NaCl, the partially negative ends of some of the water molecules attract the positive ions (cations) while the positive poles of others attract the negative ions (anions).

  9. Ionic Compounds in Water: An ionic solute is said to dissociate into ions upon dissociation. When a salt dissolves in water ionic bonds are broken. When a salt precipitates from water, ionic bonds are formed NaCl (s) + H2O (l) → Na+ (aq) + Cl– (aq)

  10. Water molecules interact favorably with both cations and anions.Water “solvates” ions. The negative pole of water interacts with cations. The positive pole of water interacts with anions.

  11. Lies my Chemistry Teacher told me: "Like dissolves like” substances with similar intermolecular attractive forces tend to be soluble in one another. Water is polar: polar molecules like fructose dissolve in water. Hexane is non-polar: non-polar molecules do not dissolve in water. But DMSO and acetone are polar. Hexane (non-polar) dissolves in DMSO (polar) and acetone (polar).

  12. Reality: Water is special. It is a very polar solvent, but that is only part of the story. Non-polar substances are driven out of water by “the hydrophobic effect”, which is a property of water, and of water only. The molecular interactions between water molecules drive non-polar solutes out of solution. The interactions between solute molecules are not the issue. Hexane does not care whether it is sitting next to other hexanes or next to water molecules. Water is unstable when it is directly adjacent to hexane.

  13. Electrolytes and Non-Electrolytes: Electrolytes: Mainly ionic compounds such as potassium sulfate (K2SO4) and sodium chloride (NaCl), which, as a solute, increase the electrical conductance of water when they dissolve. Non-Electrolytes: Do not increase the conductance of water when they dissolve. Strong electrolytes: Increase conductivity more than weak electrolytes.

  14. The Composition of Solutions The Molarity (M) of a solute in a solution is the amount of the solute (n in moles) divided by the total volume (V in liters) of the solution. Example: Obtain 180 g of fructose. Add it to a volumetric flask. Add water, with stirring, until the total volume is 1 L. What is the molarity? MWt = 6x12 + 12 + 6x16 = 180 g/mol n = 180 g (180 g/mol)-1 = 1 mol M = 1 mol/1L = 1 = mol L-1 How to make a mistake: If you add 180 g of fructose to 1 L of water, the final volume will be greater than 1 L, and the concentration will be < 1 M.

  15. Calculate the molarity (M in mol / L) of a solution obtained by dissolving 10.0 g of Al(NO3)3 in enough water to make 250.0 mL of solution. nAl(NO3)3 = mass Al(NO3)3 g x [mol weight g/mol]-1 Vsolution = given MAl(NO3)3 = M (M = mol/L

  16. Types of Chemical Reactions • Precipitation Reactions (of ionic compounds) • Ionic Equations and Net Ionic Equations • Predicting Precipitation Reactions • Acids and Bases (move protons) • Arrhenius Acids and Bases Theory • Strong and Weak acid • Naming Acids • Weak Bases • Modifying the Arrhenius model • Acid-Base Titrations • Oxidation-Reduction Reactions (move electrons) • Oxidizing and Reducing Agents • Oxidation Number • Types of Redox Reactions • Combination and Decomposition • Oxygenation • Hydrogenation • Displacement Reactions • Disproportionation

  17. Predicting Dissolution and Precipitation Reactions CHEM1310 McKelvy Lecture

  18. Ionic Equations (precipitation equation) What happens when you add solutions of barium chloride and potassium sulfate? BaCl2(aq) + K2SO4(aq) → ??? 1) Identify all the possible ions, rewrite the equation with explicit ions. Ba2+(aq) + 2 Cl-(aq) + 2 K+(aq) + SO42-(aq) → ??? 2) Look at a table of solubilities. Notice that barium sulfate is insoluble. Write the ‘complete ionic equation. Ba2+(aq) + 2 Cl-(aq) + 2 K+(aq) + SO42-(aq) → BaSO4(s) + 2 K+(aq) + 2 Cl-(aq) 3) Identify the ‘spectator ions’ (chloride and potassium). 4) Write out the “Net Ionic Equation”. BaCl2(aq) + K2SO4(aq) → BaSO4(s) + 2 KCl(aq)

  19. Write a net ionic equation to represent the formation of the precipitate observed when aqueous solutions of CaCl2 and NaF are mixed. Identify the precipitate and the spectator ions. CaCl2(aq) + NaF(aq) → ??? Rewrite with explicit ions (complete ionic equation. Ca2+(aq) + 2Cl-(aq) +Na+(aq) + F-(aq) → ??? Check the solubility table. Note that CaF2 is insoluble. Balance the equation. Ca2+(aq) + 2 Cl-(aq) + 2 Na+(aq) + 2 F-(aq) → CaF2(s) + 2 Na+(aq) + 2 Cl-(aq) Na+ and Cl- are spectator ions. CaF2 is the precipitate.

  20. Summary: Three Types of Equations Used to Describe Reactions in Solution • Molecular Equation • Complete Ionic Equation • Net Ionic Equation BaCl2(aq) + K2SO4(aq) → BaSO4(s) + 2 KCl(aq) Ba2+(aq) + 2 Cl-(aq) + 2 K+(aq) + SO42-(aq) → BaSO4(s) + 2 K+(aq) + 2 Cl-(aq) Ba2+(aq) + SO42-(aq) → BaSO4(s)

  21. Acids and Bases and Their Reactions(more in Chapter 7) Acid Base Theory • Arrhenius Acids and Bases Acids are • H+ donors • Bases are OH- donors • Arrhenius Broadened Definition • Acids increase H+ concentration or [H+] increases • Bases increase OH- concentration or [OH-] increases • Brønsted-Lowry Acids and Bases (1923) • Acids donate H+ • Bases accept H+ Arrhenius 1903 Nobel Prize

  22. Acids and Bases Neutralization reaction: HCl(aq) + NaOH(aq) → H2O(l) + NaCl(aq) H+(aq) + Cl-(aq) + Na+(aq) + OH-(aq) → H2O(l) + Na+(aq) + Cl-(aq) HCl (the acid) has donated a proton (H+) to OH- OH- (the base) has accepted a proton from HCl Neutralization occurs at the equivalence point, where nbase = nacid number of moles of acid = number of moles of base

  23. The titration of an acid with a base. End point = Equivalence point = Stoichiometric point Base (Titrant): Standard Solution Indicator

  24. Example Equivalence Point Calculation Compute the molarity of a solution of sodium hydroxide if 26 ml of solution must be added to a solution containing 0.53 g of KH5C8O4 (KHP, potassium hydrogen phthalate) to reach the end-point. If you see “end-point” or “equivalence point”, think nacid = nbase 1) nacid = 0.53 g (39 + 5 + 96 + 76 g/mol)-1 =2.5 mmol 2) nbase = 2.5 mmol NaOH [NaOH] = 2.5 mmol / 26 mL [x] is molarity of x = 0.10 mmol/ml = 0.10 mol/L = 0.10 M

  25. Arrhenius Acids and Bases: Strong acids: Hydrochloric acid (HCl) Sulfuric acid (H2SO4) Perchloric acid (HClO4) Nitric acid (HNO3) Hydroiodic acid (HI) Hydrobromic acid (HBr) Chloric acid (HClO3) Strong bases: Sodium hydroxide (NaOH) Potassium hydroxide (KOH) Strong acids and bases are strong electrolytes. Strong acids and bases dissociate to ions completely in water.

  26. Weak Acids Acetic acid (CH3COOH) Phosphoric acid (H3PO4) Formic acid (HCOOH) Hydrofluoric acid (HF) Oxalic acid (H2C2O4) CH3COOH(aq) → CH3COO−(aq) + H+(aq) (partial dissociation of acetic acid) Neutralization Reaction (weak or strong acid plus strong base) CH3COOH(aq) + NaOH(aq) → CH3COO− (aq) + H2O(l) (molecular eq.) CH3COOH(aq) + Na+(aq) + OH-(aq) → Na+(aq) + CH3COO−(aq) + H2O(l) (complete ionic eq.) Weak acids partially dissociate in water. Some on the acid(aq) remains. Weak acids are weak electrolytes.

  27. Weak Bases Ammonia (NH3) acetate CH3COO− NH3(aq) + H2O(l)→ NH4+(aq) + OH-(aq) CH3COO-(aq) + H2O(l)→ CH3COOH(aq) + OH-(aq) Weak acids partially dissociate in water. Weak base are weak electrolytes.

  28. Oxidation-Reduction Reactions Redox: an extensive and important class of reactions that is characterized by the transfer of electrons. 2 Mg(s) + O2(g) → 2 MgO(s) Magnesium is oxidized: it gives up electrons as the oxidation state of its atoms increases from zero to +2. Oxygen is reduced: it gains electrons as the oxidation state decreases from zero to -2 (i.e., becomes more negative).

  29. Oxidation States: Examples • NaCl • CrO3 • TlCl3 • Mn3N2 • O2 • H2 • CH4 • CO2 • C6H12O6 • MnO4-

  30. Oxidation-Reduction reactions → O2 + H2 H2O → Cl2 + Na NaCl → H+ + Mg Mg++ + H2

  31. A summary of an oxidation-reduction process, in which M is oxidized and X is reduced.

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