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Solutions!. What is a solution? A homogeneous mixture! Made up of a solute and solvent.

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Presentation on theme: "Solutions!. What is a solution? A homogeneous mixture! Made up of a solute and solvent."— Presentation transcript:

1 Solutions!

2 What is a solution? A homogeneous mixture! Made up of a solute and solvent.

3 2 Parts of a Solution Solvent – Does the dissolving. Usually present in the larger amount. Solute – Is dissolved. Usually present in the smaller amount.

4 Aqueous Solutions Water is a good solvent because the molecules are polar. The oxygen atoms have a partial negative charge. The hydrogen atoms have a partial positive charge.

5 Hydration The process of breaking the ions of salts apart. Ions have charges and attract the opposite charges on the water molecules.

6 Hydration H H O H H O H H O H H O H H O H H O H H O H H O H H O

7 NaCl  Na + + Cl - Al(NO 3 ) 3  Al 3+ + 3NO 3 - MgCl 2  Mg 2+ + 2Cl -

8 Soluble Vs. Insoluble Soluble: when attraction between ions and water molecules stronger than the attraction between ions Insoluble: when attraction between the ions is stronger than the attraction between the ions and water molecules

9 Electrolytes Conducts an electric current when placed in aqueous solution or in molten state. (Soluble) Non-electrolytes do not conduct an electric current in either state. (Insoluble)

10 Properties of Solutions Solubility – how much dissovles Determined by a few factors. - Nature of solvent and solute. - Agitation. - Surface area of solute. - Temperature.

11 Nature of Solvent and Solute “Like dissolves like.” Water is polar. It has small charges. It dissolves ionic and polar compounds. Non-polar will not dissolve.

12 Agitation Shaking increases the interaction between water and the solute. What do you do when you make Kool- Aid?

13 Surface Area More exposed surface area means the water molecules can interact more with the solute. A powder dissolves better than a solid.

14 Solubility of Solids Increases with temp

15 Solubility of Gases - Decreases with temp inc.

16 Concentration A quantitative measure of the amount of solute dissolved in a given quantity of solution Can use a few terms to define concentration. Solubility Molarity Molality

17 Solubility The amount of a substance that dissolves in a given quantity of a solvent at a given temperature. (g solute/100g solvent) Saturated vs. Unsaturated vs. Supersaturated

18 Solubility Saturated Solvent can’t dissolve more solute Unsaturated Solvent can dissolve more solute Supersaturated Solvent dissolved more than expected

19

20 Molarity M = moles solute/liters of solution Uses the symbol M Example: 2 moles of glucose are dissolved to make 2 liters of solution. M = 2 moles/2L = 1M

21 Sample Problem #1 A saline solution contains 0.90g of NaCl in exactly 100mL of solution. What is the molarity of the solution?

22 Sample Problem #2 How many moles are present in 250mL of 0.24M calcium chloride?

23 Molality Another term used to describe the concentration. Uses the symbol m. m = moles solute/kg solvent

24 Sample Problem #1 What is the resulting molality when 30g of NaCl is dissolved in 100g of water?

25 Dilutions Use a concentrated stock solution to make a more dilute solution. M 1 V 1 = M 2 V 2

26 Dilutions We want 4.5L of a 1M HCl solution. How much 12M HCl should we use?

27 Colligative Properties Depend on the number of particles dissolved in a given mass of solvent. Freezing point depression Boiling point elevation Vapor pressure lowering

28 Colligative Properties Freezing point depression and boiling point elevation both use a similar equation. ∆T = K(m)(i) K is a constant based on the solvent m is the molality of the solution i is the van’t Hoff factor, the number of particles a solute creates in solution

29 Van’t Hoff Factor NaCl Al 2 (CO 3 ) 3 MgCl 2

30 Colligative Properties Which of the following will lower the freezing point of water the most? a. NaCl b. Glucose c. CaI 2 d. Al 2 (CO 3 ) 3

31 Types of Reactions  Precipitation reactions When aqueous solutions of ionic compounds are poured together a solid forms. A solid that forms from mixed solutions is a precipitate If you’re not a part of the solution, you’re part of the precipitate

32 Precipitation reactions NaOH(aq) + FeCl 3 (aq)   NaCl(aq) + Fe(OH) 3 (s) is really Na + (aq)+OH - (aq) + Fe +3 (aq) + Cl - (aq)   Na + (aq) + Cl - (aq) +Fe(OH) 3 (s) So all that really happens is OH - (aq) + Fe +3 (aq)  Fe(OH) 3 (s) Double replacement reaction

33 Precipitations Reactions Only happen if one of the products is insoluble Otherwise all the ions stay in solution- nothing has happened. Need to memorize the rules for solubility

34 Solubility Rules  All nitrates are soluble  Alkali metals ions and NH 4 + ions are soluble  Halides are soluble except Ag +, Pb +2, and Hg 2 +2  Most sulfates are soluble, except Pb +2, Ba +2, Hg +2,and Ca +2

35 Solubility Rules  Most hydroxides are insoluble except NaOH and KOH  Sulfides, carbonates, chromates, and phosphates are insoluble  Lower number rules supersede so Na 2 S is soluble

36 Precipitation reaction We can predict the products Can only be certain by experimenting AgNO 3 (aq) + KCl(aq)  Zn(NO 3 ) 2 (aq) + BaCr 2 O 7 (aq) 

37 Three Types of Equations Molecular Equation- written as whole formulas, not the ions. K 2 CrO 4 (aq) + Ba(NO 3 ) 2 (aq)  Complete Ionic equation show dissolved electrolytes as the ions. 2K + + CrO 4 -2 + Ba +2 + 2 NO 3 -  BaCrO 4 (s) + 2K + + 2 NO 3 - Spectator ions are those that don’t react.

38 Three Type of Equations Net Ionic equations show only those ions that react, not the spectator ions Ba +2 + CrO 4 -2  BaCrO 4 (s) Write the three types of equations for the reactions when these solutions are mixed. iron (III) sulfate and potassium sulfide


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