States Beyond Gasses?

1. States Beyond Gasses? • Why do other states exist? • Attraction! • Let’s look at the ways molecules attraction one another:

2. Intermolecular Forces The attractions between molecules are not nearly as strong as the intramolecular attractions that hold compounds together.

3. Intermolecular Forces They are, however, strong enough to control physical properties such as boiling and melting points, vapor pressures, and viscosities.

4. Intermolecular Forces These intermolecular forces as a group are referred to as van der Waals forces.

5. van der Waals Forces • Dipole-dipole interactions • Hydrogen bonding • London dispersion forces

6. Dipole-Dipole Interactions • Molecules that have permanent dipoles are attracted to each other. • The positive end of one is attracted to the negative end of the other and vice-versa. • These forces are only important when the molecules are close to each other.

7. Dipole-Dipole Interactions The more polar the molecule, the higher is its boiling point.

8. London Dispersion Forces While the electrons in the 1s orbital of helium would repel each other (and, therefore, tend to stay far away from each other), it does happen that they occasionally wind up on the same side of the atom.

9. London Dispersion Forces At that instant, then, the helium atom is polar, with an excess of electrons on the left side and a shortage on the right side.

10. London Dispersion Forces Another helium nearby, then, would have a dipole induced in it, as the electrons on the left side of helium atom 2 repel the electrons in the cloud on helium atom 1.

11. London Dispersion Forces London dispersion forces, or dispersion forces, are attractions between an instantaneous dipole and an induced dipole.

12. Factors Affecting London Forces • The strength of dispersion forces tends to increase with increased molecular weight. • Larger atoms have larger electron clouds which are easier to polarize.

13. Hydrogen Bonding • The dipole-dipole interactions experienced when H is bonded to N, O, or F are unusually strong. • We call these interactions hydrogen bonds.

14. Hydrogen Bonding • Hydrogen bonding arises in part from the high electronegativity of nitrogen, oxygen, and fluorine. Also, when hydrogen is bonded to one of those very electronegative elements, the hydrogen nucleus is exposed.

15. Solutions • A solution is a homogeneous mixture. • A solution is composed of a solute dissolved in a solvent. • Solutions exist in all three physical states:

16. Polar Molecules • When two liquids make a solution, the solute is the lesser quantity, and the solvent is the greater quantity. • Recall that a net dipole is present in a polar molecule. • Water is a polar molecule. Chapter 14

17. Polar and Nonpolar Solvents • A liquid composed of polar molecules is a polar solvent. Water and ethanol are polar solvents. • A liquid composed of nonpolar molecules is a nonpolar solvent. Hexane is a nonpolar solvent. Chapter 14

18. Like Dissolves Like • Polar solvents dissolve in one another. • Nonpolar solvents dissolve in one another. • This is the like dissolves like rule. • Methanol dissolves in water, but hexane does not dissolve in water. • Hexane dissolves in toluene, but water does not dissolve in toluene.

19. Miscible and Immiscible • Two liquids that completely dissolve in each other are miscible liquids. • Two liquids that are not miscible in each other are immiscible liquids. • Polar water and nonpolar oil are immiscible liquids and do not mix to form a solution.

20. Solids in Solution • When a solid substance dissolves in a liquid, the solute particles are attracted to the solvent particles. • When a solution forms, the solute particles are more strongly attracted to the solvent particles than other solute particles. • We can also predict whether a solid will dissolve in a liquid by applying the like dissolves like rule.

21. Like Dissolves Like for Solids • Ionic compounds, like sodium chloride, are soluble in polar solvents and insoluble in nonpolar solvents. • Polar compounds, like table sugar (C12H22O11), are soluble in polar solvents and insoluble in nonpolar solvents. • Nonpolar compounds, like naphthalene (C10H8), are soluble in nonpolar solvents and insoluble in polar solvents.

22. The Dissolving Process • When a soluble crystal is placed into a solvent, it begins to dissolve. • When a sugar crystal is placed in water, the water molecules attack the crystal and begin pulling part of it away and into solution. • The sugar molecules are held within a cluster of water molecules called a solvent cage.

23. Dissolving of Ionic Compounds • When a sodium chloride crystal is placed in water, the water molecules attack the edge of the crystal. • In an ionic compound, the water molecules pull individual ions off of the crystal. • The anions are surrounded by the positively charged hydrogens on water. • The cations are surrounded by the negatively charged oxygen on water.

24. Rate of Dissolving • There are three ways we can speed up the rate of dissolving for a solid compound: • Heating the solution • This increases the kinetic energy of the solvent, and the solute is attacked faster by the solvent molecules. • Stirring the solution • This increases the interaction between solvent and solute molecules. • Grinding the solid solute • There is more surface area for the solvent to attack.

25. Solubility of Solids and Temperature • The solubility of a compound is the maximum amount of solute that can dissolve in 100 g of water at a given temperature. • In general, a compound becomes more soluble as the temperature increases.

26. Saturated Solutions • A solution containing exactly the maximum amount of solute at a given temperature is a saturated solution. • A solution that contains less than the maximum amount of solute is an unsaturated solution. • Under certain conditions, it is possible to exceed the maximum solubility of a compound. A solution with greater than the maximum amount of solute is a supersaturated solution.

27. Supersaturated Solutions • At 55 C, the solubility of NaC2H3O2 is 100 g per 100 g water. • If a saturated solution at 55 C is cooled to 20 C, the solution is supersaturated. • Supersaturated solutions are unstable. The excess solute can readily be precipitated.

28. Supersaturation • A single crystal of sodium acetate added to a supersaturated solution of sodium acetate in water causes the excess solute to rapidly crystallize from the solution.

29. Concentration of Solutions • The concentration of a solution tells us how much solute is dissolved in a given quantity of solution. • We often hear imprecise terms such as a “dilute solution” or a “concentrated solution.” • There are three precise ways (that we will talk about) to express the concentration of a solution: • Mass/mass percent • Molarity

30. mass of solute g solute x 100% = m/m % mass of solution x 100% = m/m % g solute + g solvent Mass Percent Concentration • Mass percent concentration compares the mass of solute to the mass of solvent. • The mass/mass percent (m/m %) concentration is the mass of solute dissolved in 100 g of solution.

31. 5.50 g NaCl x 100% = m/m % 5.00 g NaCl + 97.0 g H2O 5.00 g NaCl x 100% = 4.90 % 102 g solution Calculating Mass/Mass Percent • A student prepares a solution from 5.00 g NaCl dissolved in 97.0 g of water. What is the concentration in m/m %?

32. 100 g solution 25.0 g dextrose x 5.00 g dextrose = 500 g solution Mass Percent Calculation • What mass of a 5.00 m/m % solution of dextrose contains 25.0 grams of dextrose? • We want grams solution; we have grams dextrose.

33. moles of solute = M liters of solution Molar Concentration • The molar concentration, or molarity (M), is the number of moles of solute per liter of solution, and is expressed as moles/liter. • Molarity is the most commonly used unit of concentration.

34. 1 mol NaOH 24.0 g NaOH x = 6.00 M NaOH 40.00 g NaOH 0.100 L solution Calculating Molarity • What is the molarity of a solution containing 24.0 g of NaOH in 0.100 L of solution? • We also need to convert grams NaOH to moles NaOH (M = 40.00 g/mol).

35. 0.100 mol K2Cr2O7 294.2 g K2Cr2O7 x 250.0 mL solution x 1000 mL solution 1 mol K2Cr2O7 Molar Concentration Problem • How many grams of K2Cr2O7 are in 250.0 mL of 0.100 M K2Cr2O7? • We want mass K2Cr2O7; we have mL solution. = 7.36 g K2Cr2O7

36. 1 mol HCl 1000 mL solution 9.15 g HClx x 36.46 g HCl 12.0 mol HCl Molar Concentration Problem, Continued • What volume of 12.0 MHCl contains 9.15 g of HCl solute (M = 36.46 g/mol)? • We want volume; we have grams HCl. = 20.9 mL solution

37. moles of solute = m kilogram of solvent Molality • Molarity (m), is the number of moles of solute per kilogram of solvent, and is expressed as moles/kg. • Molality is the 2nd most commonly used unit of concentration. • Very important when we talk about “Colligative Properties”

38. 187.77 g AgBr 0.100 mol AlBr3 3 mol AgBr 37.5 mL solnx x x 1 mol AgBr 1000 mL soln 1 mol AlBr3 Solution Stoichiometry Problem • What mass of silver bromide is produced from the reaction of 37.5 mL of 0.100 M aluminum bromide with excess silver nitrate solution? AlBr3(aq) + 3 AgNO3(aq) → 3 AgBr(s) + Al(NO3)3(aq) • We want g AgBr; we have volume of AlBr3. = 2.11 g AgBr

39. Colligative Properties • Changes in colligative properties depend only on the number of solute particles present, not on the identity of the solute particles. • Among colligative properties are • Boiling point elevation • Freezing point depression

40. Boiling Point Elevation • The change in boiling point is proportional to the molality of the solution: Tb = Kb  m where Kb is the molal boiling point elevation constant, a property of the solvent. Tb is added to the normal boiling point of the solvent.

41. Boiling Point Elevation • The change in freezing point can be found similarly: Tf = Kf  m • Here Kf is the molal freezing point depression constant of the solvent. Tf is subtracted from the normal boiling point of the solvent.

42. Note that in both equations, T does not depend on what the solute is, but only on how many particles are dissolved. Tb = Kb  m Tf = Kf  m Boiling Point Elevation and Freezing Point Depression

43. Colligative Properties of Electrolytes Since these properties depend on the number of particles dissolved, solutions of electrolytes (which dissociate in solution) should show greater changes than those of nonelectrolytes.

44. van’t Hoff Factor One mole of NaCl in water gives to two moles of ions.

45. van’t Hoff Factor • We modify the previous equations by multiplying by the van’t Hoff factor, i. Tf = Kf m  i