classification of matter:

1. Chapter 5: Solutions 5.1 Classifying Solutions A. The Basics • classification of matter: All Matter Mixtures Pure Substances Heterogeneous (Mechanical Mixtures and Colloids) Homogeneous (Solutions) Elements Compounds

2. is any that has matter solid, liquid or gas mass and volume • are a type of matter that has pure substances definite fixed composition eg) elements and compounds • are combinations of matter that can be mixtures do NOT have definite proportions separated by physical means and

3. have and the different components are heterogeneousmixtures (mechanical mixtures) varying composition usually visible • are which have solutions homogeneous mixtures uniform composition and the different components are not visible • composed of at least one substance dissolved in another

4. solvent • is the and is usually the substance present in the dissolver, largest quantity (mass, volume, amount) eg) water • is what is solute dissolved in the solvent eg) salt

5. B. Types of Solutions • both solvents and solutes can be solids, liquids or gases • you can have various combinations of solute and solvent phases eg) liquid in liquid – ethylene glycol (antifreeze) solid in liquid – Kool Aid gas in liquid – carbonated beverages solid in solid – alloys liquid in solid – mercury amalgam fillings

6. an solution is any solution in which is the aqueous water solvent • water is called the therefore we will be concerned mainly with “universal solvent” aqueous solutions • water dissolves a lot of different solutes because of it’s …this is good because …also bad because dissolve in it easily unique properties 75% of the earth (and our bodies) is water toxins

7. C. Dissolving and the Forces of Attraction • is a dissolving physical change • the of a solid solute are held together by molecules or ions bonds • when dissolving occurs, these bonds and the of the solute become break ions or molecules attracted to the solvent particles H2O NaCl

8. the 3 processes involved in dissolving are: 1. bonds between molecules or ions of broken endothermic (requires energy) solute – always 2. bonds between molecules of broken solvent – always endothermic 3. bonds between molecules or ions of form solute and solvent – always exothermic

9. overall energy change in dissolving is equal to the of the three steps sum (law of conservation of energy) • if, the overall dissolving process is more energy is released than is required exothermic • if, the overall dissolving process is less energy is released than is required endothermic

10. D. Electrolytes vs Non-Electrolytes • are aqueous solutions that electrolytes (weak and strong) conduct electricity eg) all soluble ionic compounds, very polar molecular compounds (like acids, ammonia) • are aqueous solutions that non-electrolytes do not conductelectricity eg) molecular compounds in solution

11. occurs when when they are dissolved in an dissociation ionic compounds break apart into their ions aqueous solution • are used to show what happens to a substance when it is put into dissociation equations water • you can have 4 situations: 1. insoluble ionic or molecular compounds • they to any great extent do not dissolve solubility table • use the for ionic compounds  eg) AgCl(s) AgCl(s)  C25H52(s) C25H52(s)

12. 2. soluble ionic compounds • dissolve to a great extent to form ions in solution • are broken ionic bonds • use solubility table • includes which turn litmus paper bases blue • the number of ions balance eg) NaCl(s)  Na+(aq) + Cl(aq) Ca(OH)2(s)  Ca2+(aq) + 2 OH(aq)

13. 3. soluble molecular compounds • dissolve to form molecules in solution intermolecular forces (LD, DD, HB) • are broken  eg) C12H22O11(s) C12H22O11(aq)

14. 4. acids • they are but are molecular compounds very polar • dissolve to form in solution ions (ionize) • turns litmus paper red • the number of ions balance H3PO4(s)  eg) 3H+(aq) + PO43(aq)

15. Examples Write the equations to show what happens to each of the following in water: 1. potassium chloride KCl(s)  K+(aq) + Cl-(aq) 2. carbon dioxide  CO2(g) CO2(aq)

16. 6. gasoline  C8H18(l) C8H18(l) 7. barium sulphate BaSO4(s)  BaSO4(s)

17. 3. solid hydrogen nitrate  HNO3(s) + H+(aq) NO3(aq) 4. aluminum sulphate Al2(SO4)3 (s)  Al3+ (aq) 2 + 3 SO42- (aq) 5. sodium phosphate decahydrate 3 Na3PO4  10H2O(s)  Na+ (aq) + PO43- (aq)

18. many chemical reactions (either in our bodies or in the lab) are if the reactants are not dissolvedin very slow water • dissolving in water allows the solute particles to separate, disperse and collide with other solute particles • particle collisions are necessary for to occur reactions

19. 5.2 Solubility A. The Definitions amount of a solute that dissolves in a given quantity of solvent solubility • the of a solute is the at a given temperature eg) has a solubility of NaCl(s) 36 g/100mL at 20C • an is a solution that does have the maximum amount of solute dissolved in it unsaturated solution not

20. saturated solution • a is a solution that contains the maximum amount of a dissolved solute at a given temperature • some will be present undissolved solute • the solution may still be able to dissolveother solutes

21. a contains more dissolved solute than its solubility supersaturated solution at a given temperature sodium acetate video

22. B. Range of Solubility • the degree to which a solute is soluble depends on the between: strength of attraction 1. the solute particles 2. the solute particlesand the solvent particles • solutes can be: less than 0.1 g/100mL, 1. insoluble – although there is still a tiny bit of dissolving 2. slightly soluble – between 0.1 g/100 mL and 1 g/100 mL 3. soluble – greater than 1 g/100 mL

23. note that these general rules for solubility do not apply to …the numbers for gases are much and still considered gases lower soluble eg) has a solubility of yet this is considered O2(g) 0.009 g/100 mL at 20C soluble

24. C. A Closer Look dissolves • once a solute , it appears that the solute-solvent bonds don’t break • when you study a saturated solution, the amount (mass or moles) of undissolved solute at the bottom remains unchanged • over time, undissolved particles become dissolved and dissolved particles crystallize • a saturated solution is said to be in a state of equilibrium

25. dissolving crystallization (dissolving) • equilibrium occurs when a process and the reverse processtake place at the (crystallization) same rate

26. CuSO4(s)  Cu2+(aq) + SO42(aq) is Cu2+(aq) + SO42(aq)  CuSO4(s) is What we do is use a double arrow to show equilibrium: CuSO4(s) Cu2+(aq) + SO42(aq) eg) dissolving crystallization ⇌

27. D. Temperature and Solubility • the for a substance must be provided with a certain and it depends on the of the solute solubility value temperature state Solids • when a dissolves in a , the holding the solid together must be solid liquid bonds broken • at , the particles of the solute and solvent have more higher temperatures energy • as temperature the solubility of a solid increases, increases

28. Liquids • the particles of a liquid are held together asas the particles in a solid not strongly • when a dissolves in a , additional is needed liquid liquid energy not • the solubility of most liquids is not affected by temperature

29. Gases • gas particles move and have a great deal of quickly energy • when a dissolves in a , the particles must gas liquid lose energy • the temperature would the gas particles and make it for them to dissolve increasing speed up harder • as temperature the solubility of a gas increases, decreases

30. E. Pressure and Solubility • a change in has a negligible effect pressure on the solubility of solids or liquids • the solubility of are greatly affected by gases pressure • as pressurethe solubility of a gas increases, increases eg) pop bottles, cracking knuckles, the “bends”

32. 5.3 Concentration A. Solutions in our Society • solutions are all around us and affect our lives in many ways • we have developed many to meet the and needs of humans technologies personal industrial eg) hair products, breathalyzer test, tests to monitor drinking water

33. in using solutions in our lives, care must be taken to ensure responsible use eg) – set up in Canada in 1985 to ensure that companies that deal with chemicals are using them in a safe and appropriate manner Responsible Care Program

34. using solutions can have intended and unintended consequences for humans and the environment • released into the environment are taken in by organisms toxins • this can have toxic effects immediate eg) H2S(g)…sour gas will cause a loss of consciousness at 700 ppm

35. sometimes the of the food chain are by the toxins since the levels are not that high…but the toxins are then passed along to the of the food chain lower levels unaffected upper levels • organisms at the of the food chains (like humans!) can end up with of these chemicals top toxic levels eg) mercury, lead, arsenic, PCB’s, DDT etc.

36. is the concentration of toxins as you biomagnification (bioaccumulation) move up the food chain

37. we must assess the andof using technologies that contain certain substances risks benefits eg) heavy metals released from mineral processing, power plants – mercury, arsenic

38. B. Ways of Expressing Concentration the amount of solute relative to the amount of solvent • concentration is dilute = solute,solvent low high concentrated = high solute, solvent low • concentration can be expressed in a variety of ways: ppm, ppb, ppt, % mass, % volume, mol/L, mmol/L, mg/dL, %

39. these expressions can be seen on many products that we use in our daily lives and in industry eg) vinegar – toothpaste – blood cholesterol levels – breathalyzer - % (g of alcohol per 100 mL of blood) eg. 5% acetic acid by volume 0.243% w/w NaF 250 mg/dL 0.08

40. 1. Concentration as Percent by Mass % mass (m/m) = mass solute (g) __ × 100 mass solution (solute + solvent) (g) Example An aqueous solution with a mass of 55.5 g contains 2.85 g of solute dissolved in it. What is the percent by mass of the solute in the solution? 2.85 g percent mass (m/m) = × 100 55.5 g = 5.14 %

41. 2. Concentration in Parts per Million ppm = parts per million ppb = parts per billion mass solute (g) __ ppm = ×106 ppm mass solution (g)

42. Example Carbon monoxide is deadly at 900 ppm. If there is 8.50 g of carbon monoxide in a room that contains 11.5 kg of air, what is the concentration of CO(g) in ppm? Is the level of carbon monoxide deadly? ppm = mass solute (g) ×106 ppm mass solution (g) = 8.50 g ×106 ppm 11 500 g = 739 ppm no, it’s not deadly

43. 3. Molar Concentration molarity = number of moles of solute per litre of solvent c = n v n = m M where: c = concentration in mol/L n = number of moles in mol v = volume in L m = mass in g M = molar mass in g/mol

44. Example 1 A sample of 0.900 mol of NaCl is dissolved to give 0.500 L of solution. What is the molar concentration of the solution? n = 0.900 mol v = 0.500 L c = n v = 0.900 mol 0.500 L = 1.80 mol/L

45. Example 2 Calculate the concentration of 100 mL of a solution containing 0.300 mol of sulphuric acid. n = 0.300 mol v = 0.100 L c = n v = 0.300 mol 0.100 L = 3.00 mol/L

46. Example 3 Calculate the molar concentration of a 250 mL solution that has 3.2 g of NaCl dissolved in it. m = 3.2 g v = 0.250 L M = 58.44 g/mol n = m M = 3.2 g 58.44 g/mol = 0.0547…mol c = n v = 0.0547… mol 0.250 L = 0.22 mol/L

47. Example 4 Calculate the number of moles of Pb(NO3)2 needed to make 500 mL of a 1.25 mol/L solution. c = 1.25 mol/L v= 0.500 L n = cv = (1.25 mol/L)(0.500 L) = 0.625 mol

48. Example 5 How many litres of 4.22 mol/L solution would contain 3.69 mol of BaCl2? c = 4.22 mol/L n = 3.69 mol v = n c = 3.69 mol 4.22 mol/L = 0.874 L

49. Example 6 Calculate the mass of the salt required to prepare 1.50 L of a 0.565 mol/L solution of K3PO4. n = cv = (0.565 mol/L)(1.50 L) = 0.8475 mol c = 0.565 mol/L v = 1.50 L M = 212.27 g/mol m = nM = (0.8475 mol)(212.27 g/mol) = 180 g

50. D. Molar Concentration of Ions • you may have to calculate the concentration ofin solution ions • once you have your dissociation equation, you can calculate the concentration of ions in solution using the mole ratio wanted given • are anions negative ions • are cations positive ions