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  1. Reactions in Solution • Some Important Definitions • Electrolytes • Reactions in Solution • Ionic Equations • Single Replacement Reactions • Concentration • Solution Stoichiometry • Titrations

  2. Solutions • A solutionA homogeneous mixture of two • or more components. Sugar in water Oxygen in water Air Dental fillings Saline

  3. A solution • In a solution • The solute can’t be filtered out. • The solute always stays mixed. • Particles are always in motion. • Volumes may not be additive. • A solution will have different properties than the solvent • A solution consists of two component types. • solvent - component in the greater extent • solute - component in the lesser extent • (You may have more than one.)

  4. Physical states of solutions • Solutions can be made that exist in any of the three states. • Solid solutions • dental fillings, 14K gold, sterling silver • Liquid solutions • saline, vodka, vinegar, sugar water • Gas solutions • the atmosphere, anesthesia gases

  5. Solubility • A measure of how much of a solute can be dissolved in a solvent. • Common unit • - grams / 100 ml • Factors affecting solubility • Temperature • Pressure • Polarity

  6. Solubility of some substances • Temperature Solubility • Substance oC g/100 ml water • NaCl (s) 100 39.12 • PbCl2 (s) 100 3.34 • AgCl (s) 100 0.0021 • CH3CH2OH (l) 0 - 100 infinity • CH3CH2OCH2CH3 (l)15 8.43 • O2 (g) 60 0.0023 • CO2 (g) 40 0.097 • SO2 (g)40 5.41

  7. Saturation • When a solution contains as much solute as it can at a given temperature. • Unsaturated Can still dissolve more. • Saturated Have dissolved all you can. • Supersaturated Temporarily have dissolved • too much. • Precipitate Excess solute that falls out • of solution.

  8. Properties ofaqueous solutions • There are two general classes of solutes. • Electrolytic • ionic compounds in polar solvents • dissociate in solution to make ions • conduct electricity • may be strong (100% dissociation) or weak (less than 100%) • Nonelectrolytic • do not conduct electricity • solute is dispersed but does not dissociate

  9. Dissolving ionic compounds • When an ionic solid dissolves in water, the solvent removes ions from the crystal.

  10. Dissolving ionic compounds

  11. Dissolving covalent compounds • Covalent compounds do not dissociate.

  12. Acids, bases and salts • Three types of compounds are electrolytes: • Acid - a compound that increases the concentration of hydrogen ions in water. • HCl H+ + Cl- • Base - a compound that increases the concentration of hydroxide ion in water. • NaOH Na+ + OH- • Salt - the ions that remain after an acid and base react with each other - neutralization. • HCl(aq) + NaOH(aq) NaCl(aq) + H2O(l) water water water

  13. Some acids, bases and their salts • Acid Sodium salt • Name Formula Name Formula • Acetic acid HC2H3O2 Sodium acetate NaC2H3O2 • Hydrogen chloride HCl Sodium chloride NaCl • Nitric acid HNO3 Sodium nitrate NaNO3 • Phosphoric acid H3PO4 Sodium phosphate Na3PO4 • Sulfuric acid H2SO4 Sodium sulfate Na2SO4 • Base Chloride salt • Name Formula Name Formula • Sodium hydroxide NaOH Sodium chloride NaCl • Barium oxide BaO Barium chloride BaCl2 • Sodium oxide Na2O Sodium chloride NaCl • Ammonia NH3 Ammonium chloride NH4Cl

  14. Reactions between ions in solution • Neutralization is an example of a reaction between ions in solution. • When ions react, we might observe the formation of a precipitate or a gas. • AgNO3 (aq) + NaCl (aq) AgCl (s) + NaNO3 (aq) • Na2CO3 (aq)+ 2HNO3 (aq) 2NaNO3 (aq) +H2O (l) CO2 (g) • However, not all ions will react in solution. • KNO3 (aq) + NaCl (aq) No reaction • Solubility rules can help predict reactions.

  15. Some simple solubility rules • All acids are soluble. • All Na+, K+ and NH4+ salts are soluble. • All nitrate and acetate salts are soluble. • All chlorides except AgCl and Hg2Cl2 are soluble. PbCl2 is slightly soluble. • All sulfates are soluble except PbSO4, Hg2SO4, SrSO4 and BaSO4. Ag2SO4 and CaSO4 are slightly soluble. • All sulfides are insoluble except those of the Group IA (1), IIA (2) and ammonium sulfide. • All hydroxides are insoluble except those of the group IA(1) and Ba(OH)2. Sr(OH)2 and Ca(OH)2 are slightly soluble.

  16. H2O H2O Ionic equations • When ionic substances dissolve in water, they dissociate into ions. • AgNO3 Ag++ NO3- • KClK+ + Cl- • When a reaction occurs, only some of the ions are actually involved in the reaction. • Ag++ NO3- +K+ + Cl- AgCl(s) + K+ + NO3-

  17. Ionic equations • To help make the reaction easier to see, we commonly list only the species actually involved in the reaction. • Full ionic equation • Ag++ NO3- +K+ + Cl- AgCl(s) + K+ + NO3- • Net ionic equation • Ag++Cl-AgCl(s) • NO3- and K+ are referred to as spectator ions.

  18. Single replacement reaction • Where one element displaces another in a chemical compound. • H2 + CuO Cu + H2O • In this example, hydrogen replaces copper. • This type of reaction always involves oxidation and reduction (REDOX). • Since one species is replacing another, there are no spectator ions.

  19. Single replacement reactions • If various metals are in water, we observe that some are more reactive than others. • 2Na (s) + 2H2O (l) 2NaOH(aq) + H2 (g) (fast) • Ca (s) + 2H2O(l) Ca(OH)2 (s) + H2 (g) (slow) • Mg (s) + H2O (l) no reaction • This indicates that the order of reactivity of these metals towards water is • Na > Ca > Mg • We can show the reactivity of metals towards water and acids using an activity series.

  20. Activity series of metals potassium sodium Reacts violently with cold water calcium Reacts slowly with cold water magnesium aluminum zinc chromium Reacts very slowly with steam but quite reactive in acid increasing reactivity iron nickel tin lead Reacts moderately with acid copper silver platinum gold Unreactive in acid

  21. Activity series of metals -various metals in HCl Iron Zinc Magnesium

  22. Increased reactivity C N O F P S Cl Se Br Increased reactivity I Reactivity of nonmetals

  23. Concentration of solutions • We need a way to tell how much solute is in a solution - concentration. • There are many systems - we will cover four. • Weight / volume percent • Volume / volume percent • Weight / weight percent • Molarity

  24. Weight/Volume % Mass solute Total Volume Weight/Volume % = x 100 use g and ml If 5 grams of NaCl is dissolved in water to make 200 ml of solution, what is the concentration? 5 g / 200 ml * 100 = 2.5 wt/v% Saline is a 0.9 wt/v% solution of NaCl in water.

  25. Volume/Volume % Volume Solute Total Volume x 100 Volume/Volume % = Use the same units for both If 10 ml of alcohol is dissolved in water to make 200 ml of solution, what is the concentration? 10 ml / 200 ml * 100 = 5 V/V% Alcohol in wine is measured as a V/V%.

  26. Weight/Weight % Mass Solute Total Mass Weight/Weight % = x 100 Use the same units for both If a ham contained 5 grams of fat in 200 g of ham, what is the % wt/wt? 5 g / 200g * 100 = 2.5 wt/wt% On the label, it would say 97.5 % fat free.

  27. Very low concentrations • Pollutants in air and water are typically found at very low concentrations. Two common units are used to express these trace amounts. • Parts per million - ppm • Parts per billion - ppb • Both are modifications of the % system which could be viewed as parts per hundred - pph. • Both mass and volume % systems are used.

  28. volume solute volume solution volume solute volume solution Low concentrations in air • Trace amounts in are are expressed as volume/volume ratios. • ppm = x 106 • ppb = x 109 • Example. One cm3 of SO2 in one m3 of air would be expressed as 1 ppm or 1000 ppb.

  29. mass solute mass solution mass solute mass solution Low concentrations in water • Mass percentages are used for water pollutants. • ppm = x 106 • ppb = x 109 • Example. One ppm of a toxin in water is the same as 1 mg / liter since one liter of water has a mass of approximately 106 mg.

  30. moles solute mol liters of solution L M = = Molarity • Molarity • Recognizes that compounds have different • formula weights. • A 1 M solution of glucose contains the • same number of molecules as 1 M ethanol. • [ ] - special symbol which means molar • ( mol/L )

  31. Molarity • Calculate the molarity of a 2.0 L solution that contains 10 moles of NaOH. • MNaOH = 10 molNaOH / 2.0 L • = 5.0 M

  32. Molarity • What’s the molarity of a solution that has 18.23 g HCl in 2.0 liters? • First, you need the FM of HCl. • FMHCl = 1.008 x 1 H + 35.45 x 1 Cl • = 36.46 g/mol • Next, find the number of moles. • molesHCl = 18.23 gHCl / 36.46 g/mol • = 0.50 mol • Finally, divide by the volume. • MHCl = 0.50 mol / 2.0 L • = 0.25 M

  33. Solution preparation • Solutions are typically prepared by: • Dissolving the proper amount of solute and diluting to volume. • Dilution of a concentrated solution. • Lets look at an example of the calculations required to prepare known molar solutions using both approaches.

  34. Making a solution • You are assigned the task of preparing 100.0 ml of a 0.5000 M solution of sodium chloride. • What do you do? • First, you need to know how many moles of NaCl are in 100.0 ml of a 0.5 M solution. • mol = M x V (in liters) • = 0.5000 M x 0.1000 liters • = 0.05000 moles NaCl

  35. Making a solution • Next, we need to know how many grams of NaCl to weigh out. • gNaCl = mol x FMNaCl • = 0.05000 mol x 58.44 g/mol • = 2.922 grams

  36. Making a solution • Finally, you’re ready to make the solution. • Weigh out exactly 2.922 grams of dry, pure NaCl and transfer it to a volumetric flask. • Fill the flask about 1/3 of the way with pure water and gently swirl until the salt dissolves. • Now, dilute exactly to the mark, cap and mix.

  37. Dilution • Once you have a solution, it can be diluted by adding more solvent. This is also important for materials only available as solutions • M1V1 = M2V2 • 1 = initial • 2 = final • Any volume or concentration unit can be used as long as you use the same units on both sides of the equation.

  38. Dilution • What is the concentration of a solution produced by diluting 100.0 ml of 1.5 MNaOH to 2.000 liters? • M1V1 = M2V2 • M1 = 1.5 M M2 = ??? • V1 = 100.0 ml V2 = 2000 ml • M2 = M1V1 / V2 • M2 = (1.5 M) (100.0 ml) = 0.075 M • (2000. ml)

  39. Solution stoichiometry • Extension of earlier stoichiometry problems. • First step is to determine the number of moles based on solution concentration and volume. • Final step is to convert back to volume or concentration as required by the problem. • You still need a balanced equation and must use the coefficients for working the problem.

  40. Solution stoichiometry example • Determine the volume of 0.100 M HCl that must be added to completely react with 250 ml of 2.50 M NaOH • Balanced chemical equation • HCl(aq) + NaOH(aq) NaCl(aq)+ H2O (l) • The first step is to determine how many moles of NaOH we have.

  41. Solution stoichiometry example • We have 250 ml of a 2.50 M solution. • molNaOH = 0.250 L x 2.50 mol/L • = 0.625 molNaOH • From the balanced chemical equation, we know that we need one mole of HCl for each mole of NaOH. • That means we need 0.625 molHCl.

  42. ( ) 1 L 0.100 mol Solution stoichiometry example • Now we can determine what volume of our 0.100 M HCl solution is required. • L = molHCl / MHCl • = 0.625 mol • = 6.26 L

  43. Titration • Method based on measurement of volume. • You must have a solution of known concentration - standard solution. • It is added to an unknown solution while the volume is measured. • The process is continued until the end point is reached - a change that we can measure. • Acids and bases are commonly measured using titrations.

  44. Neutralization • The reaction of an acid with a base to produce a salt and water. • HCl(aq) + NaOH(aq) NaCl(aq) + H2O (l) • If we prepare a standard solution of NaOH, we can then use it to determine the concentration of HCl in a sample. • This is an example of Analytical Chemistry.

  45. Titrations Buret - volumetric glassware used for titrations. It allows you to add a known amount of your titrant to the solution you are testing. An indicator will give you the endpoint.

  46. Titrations

  47. Start End Titrations • Note the color change which indicates that the ‘endpoint’ has been reached.

  48. Indicator examples • Acid-base indicators are weak acids that undergo a color change at a known pH. pH phenolphthalein

  49. Indicator examples bromthymol blue methyl red