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Solutions Unit recALL:

Solutions Unit recALL: Mixture – Matter consisting of 2 or more substances combined physically. (can be hetero or homo geneous). Solution : Homogeneous mixtures made of very small, individual particles combined uniformly. Descriptions of Solutions….

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Solutions Unit recALL:

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  1. Solutions Unit recALL: Mixture – Matter consisting of 2 or more substances combined physically. (can be hetero or homogeneous)

  2. Solution: Homogeneous mixtures made of very small, individual particlescombined uniformly.Descriptions of Solutions… • Particles will not separate out if left alone • Cannot filter out any particles • Whole solution must be in ONE PHASE One Type of Solution: “Tincture” = Any solution that has Alcohol as the solvent.

  3. Solute = The substance that is dissolved. Ex: salt • Solvent = The substance that dissolves • the other substance. Ex: Water SEE VIDEO

  4. The most common solvent on earth is water – it is often called the Universal Solvent. • A solution with water as the solvent is called an aqueous solution.

  5. SOLVATION – Process by which solute particles separate from each other and become surrounded by solvent particles. (Occurs with molecules.) • *A particle (ion/molecule) is said to be hydratedif the solvent is water. Molecular Solid Four individual diatomic molecules clumped together as a solid Separated molecules in solution

  6. DISSOCIATION – This is a process by which ions separate • from each other and are surrounded by solvent particles. • *All acids dissociate. (Occurs with Ions.) + + Ionic solid – separate ions held together in a specific pattern. Separated ions in solution + +

  7. The rate of solvation/dissociation can be increased with 4 factors; • an increase in temperature 2) stirring or shaking 3) powdering or grinding 4) by adding more solvent.

  8. SOLVENT/SOLUTE COMBINATIONS“Like Dissolves Like”Polar Solute/Polar Solvent likely forms a solutionPolar Solute/Nonpolar Solvent likely won’t form a solutionNonpolar Solute/Polar Solvent likely won’t form a solutionNonpolar Solute/Nonpolar Solvent likely forms a solution

  9. Types of Solutions • Solutions can be in any of the 3 phases; Gas, Liquid, or Solid. • Gas example = Pure Air ! • Liquid: Miscible – 2 liquids that will mix together Immiscible – 2 liquids that will NOT mix together • Solid examples: Alloy – 2 metals uniformly mixed Amalgam – Hg combined with another metal

  10. Levels of Saturation Saturated Solution: A solution with the maximum amount of solute dissolved in the solvent under given conditions. *Examples ?? Unsaturated Solution: Solutions with less solute dissolved than is possible in the solvent under given conditions. Supersaturated Solution: A solution holding more solute dissolved in solvent than is expectedly possible under given conditions – VERY RARE!! Rate of Solution = The quantity of solute that will dissolve in a specific amount of time.

  11. Solubility The maximum quantity of solute that can be dissolved in a solvent at a specific temp. ** Like dissolves Like! FACTORS AFFECTING SOLUBILITY 1. NATURE OF SOLUTE/SOLVENT: The composition, IMF, bond strength… 2. Temperature: Increasing temperature increases Solubility, except for gases (opposite effect) 3. Pressure: No effect on solids/liquids. For gases, as you increase pressure the solubility increases.

  12. Solubility Curves Chart • This chart shows the amounts of solutes that will dissolve in 100 grams of water at specific temperatures.

  13. Dilute Solution: Small amount of solute dissolved in relation to solvent. Concentrated Solution: Large amount of solute dissolved in relation to solvent. Ex: Unsaturated - Saturated - Supersaturated – Usually dilute, but can be concentrated Concentrated Concentrated

  14. Expression Concentration • Molarity measures the concentration of a solution; • Molarity = Moles of solute / 1 Liter sol’n • (dm3 = L) • (cm3 = mL) Ex: A 3.2 “Molar” solution would have 3.2 moles of solute dissolved in 1 Liter of solution. We can also factor in how many molecules of solute too, right?! How many would be in the example above?

  15. How would you create a 2.5 M Ca(NO3)2 solution in 225 mL of water? 2.5M X 0.225 L = 0.56 moles Ca(NO3)2 0.56 moles Ca(NO3)2 are equal to 91.84 grams by conversion… So, dissolve 91.84 grams of Calcium Nitrate in 225mL of water! **How many mL would you need to make a 12M HCl solution if you had 88g HCl ?

  16. MOLALITY – A measurement of moles solute per kilogram of solvent. The symbol is m. A solution that is 1.2 m HBr, it really means that 1.2 moles of HBr is present for every 1 kg of solvent. • m = moles solute • kg solvent • NORMALITY (N) – A measurement based on the number of dissociated ions per molecule dissolved in the solution. • For every 1 molecule of HCl, one H+ ion and one Cl- ion will dissociate, making the solution 1N H+ and 1N Cl-. • But what if we were talking about CaBr2 in solution? • 1N Ca+2 and 2N Br-1. • MOLE FRACTION (X) – A measurement of moles of solute per moles of solution. It is a ratio only. • X = Moles Solute • moles soln

  17. DILUTION • A useful process and calculation to know is how to dilute solutions. - Oftentimes, the solution available may be more concentrated that you need or want. Therefore adding more solvent to a solution will dilute the concentration of the solution. - The formula for any dilution is C1V1 = C2V2. C1 is the initial concentration, V1 is the volume of the solution you will need, C2 is the desired concentration, and V2 is the total volume of the new solution that you will need. EX 1. A 12 M HCl solution is available, but 350 mL of a 2.5 M HCl soln is needed. How would you prepare the solution? • C1V1 = C2V2 • (12M) V1 = (2.5M)(350 mL) • V1 = 73 mL • To prepare the soln: Carefully measure out 73 mL of the 12 M HCl sol’n and dilute to a volume of 350 mL (add 277 mL of H2O).

  18. Colligative Properties Characteristics of a solution that are entirely dependent on the amount of solute in a solvent • Freezing Point Depression: The freezing point of a solvent decreases in temperature as you add solute. • Ex: Salt water freezes BELOW zero degrees Celsius. • Boiling Point Elevation: The boiling point of a solvent raises to a higher temperature as you add solute. • Ex: Water with a solute will boil ABOVE 100 degrees C.

  19. Electrolyte = A substance that breaks into ions and conducts electricity when in H2O. • Some are better than others! • All ionic compounds and Acids/Bases.. Nonelectrolyte = A substance whose ions will NOT conduct electricity in water. • Most covalent compounds (other than acids/bases)

  20. A reversible reaction is a reaction in which both the forward reaction and the reverse reactions can, and do occur. • Reversible reactions will eventually reach “Equilibrium”. • At equilibrium, the rate of the forward reaction is equal to the rate of the reverse reaction. • Ex: A + B  C + D • Equilibrium means the rate that A and B are forming products is equal to the rate that C and D are re-forming reactants. • EQUILIBRIUM DOES NOT MEAN THAT YOU HAVE EQUAL AMOUNTS OF A, B, C, and D!!! • Ex: Football teams on the sideline vs players on the field.

  21. Every reversible reaction has its own rate at which it occurs.. This is called the “Equilibrium Constant” (Keq). - Keq is calculated by the following: Keq = [products] - Brackets represent Concentration. [reactants] For example: 2H2 + O2 2H2O Keq = [H2O]2 The Coefficients become exponents. [H2]2[O2] Solubility Product Constant (Ksp) During Dissociation, Equilibrium in solution is reached when the solid solute is dissolving at the same rate that the dissociated ions are re-forming. - We can put this relationship into a chemical equation in the exact same manner as Keq. - The reactants, however, are left out of the equation because they are not actually present in solution.

  22. Ex: Na2SO4 Na2SO4 2Na+1 + SO4-2 Ksp = [Na+1]2 [SO4-2] What about Ca(NO3)2 ? Ksp = [Ca+2][NO3-1]2 Pg. 604: #’s 1-6

  23. Ionic Equations • First you must decide if the compound is ionic and soluble, OR molecular and soluble. • Recall that ionic compounds, acids, and bases dissociate into ions in solution. EX: KCl K+1 + Cl-1 - Molecular compounds Solvate: Ccl4 CCl4 In order to illustrate the states of matter we use: (s) (l) (g) and (aq) Soluble compounds are labeled (aq) when they are in sol’n. Step 1 = Write the formulas of the reactants Step 2 = Predict the products and balance the equation

  24. Potassium Chloride and Sodium Nitrate (in water) - H2O is left out of the equation because it does not react chemically with the solutes. KCl + NaNO3 KNO3 + NaCl Now use your Solubility Table to figure out: • Will the reactants dissociate (must be able to dissociate and soluble) • Would the products form a precipitate (a solid forming from a liquid sol’n) due to being insoluble. ** Otherwise there will be no reaction. Now separate the ions where needed and add the phases, and this is your full IONIC EQUATION: K+1(aq) + Cl-1(aq) + Na+1(aq) + NO3-1(aq) K+1(aq) + NO3-1(aq)+ Na+1(aq) + Cl-1(aq) So, the final result is No Reaction (N.R.) because no ions reacted chemically in the solution. They are presented in the same way on both sides of the equation.

  25. NET IONIC EQUATIONS You only write the reactants and products that have chemically changed. Ex: Copper I Sulfate and Magnesium Chloride Cu2SO4 + MgCl2 2 CuCl + MgSO4 2Cu+1(aq)+ SO4-2(aq)+ Mg+2(aq)+2Cl-1(aq) 2CuCl(s) + Mg+2(aq) + SO4-2(aq) So, the final “NET IONIC EQUATION” would be: 2Cu+1(aq) + 2Cl-1(aq) 2 CuCl(s) SO4-2(aq) and Mg+2(aq) are called “SPECTATOR IONS” because they don’t react chemically in the solution.

  26. GIVEN: 75.00mL of a 3.225 M Zinc Chloride soln and 100.0mL of a 2.460 M Rubidium Phosphate soln react 1) Predict the products of the reaction. (Type of reaction & chemical equation) 2) Balance the equation. 3) Write the ionic equation of the rxn. 4) Write the net ionic equation of the rxn. 5) Determine if a rxn occurred. 6) If yes, state how you would collect the product.* 7) Determine the mass of the reactants used. 8) What is the limiting reactant? 9) Determine the theoretical yield of the rxn. 10) How much excess react is left over? 11) Determine the percent yield of the rxn given that 27.32g of Zn3(PO4)2 is collected. 12) Determine the percent error.

  27. Next slides not needed for honors • End of Solutions Unit!

  28. Henry’s Law, states that the mass of a gas that can dissolve, in a given volume of a liquid solvent at a specified temperature, varies directly with the partial pressure of the gas. • That is to say that as the partial pressure of the gas increases above a liquid in a container, so will the solubility of that gas. • Ex: Soda can • When any solute dissolves into a solvent, there is an energy change that occurs. • This is energy change is called Enthalpy of Solution, Hsol. • Some substances require energy in order to allow them to dissolve. This would be an endothermic process and has a positive Hsol value. • Most solid solutes dissolving in a liquid are endothermic processes. • Some substances will release energy in order to allow them to dissolve. This would be an exothermic process and has a negative Hsol value. (typically gases)

  29. PERCENT BY MASS – A measurement of the mass of solute per the mass of solution. The Percent by Mass is a percentage only. It is often expressed as %w/w or %m/m. • % m/m = mass solute x 100 • mass soln • PERCENT BY VOLUME – A measurement of the volume of solute per the volume of solution. The Percent by Volume is a percentage only. It is often expressed as %V/V. • % V/V = volume solute x 100 • volume soln • PERCENT MASS BY VOLUME – A measurement of the mass of solute per the volume of solution. The Percent Mass by Volume is a percentage only. It is often expressed as %w/V or %m/V. • % m/V= mass solute x 100 • volume soln

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