Chapter 15 Notes Solution Chemistry PDF

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Chapter 15 Notes – Solution Chemistry Overall Summary of Chapter: Solution chemistry has to do with what happens to compounds when they are placed in water. Ionic compounds may dissolve and dissociate into individual ions (like NaCl becomes Na+ and Cl- ) or they might not (like PbS which stays PbS in water). You have a table in your book on page 920 - Table C-10 which will tell you if an ionic salt is soluble or not. Covalent compounds (like sugar – C12H22O11) do not dissociate in water. They dissolve into individual molecules of C12H22O11, but they do not become individual ions. The only exception to covalent compounds are the acids (HCl, CH3COOH…etc). these covalent compounds do separate into ions (HCl becomes H+ and Cl- . CH3COOH becomes CH3COO- and H+ . They always have a hydrogen cation (H+ ) and the “other” ion which is an anion (negative ions are called anions, positive ions are called cations). Whenever an H+ is present in water, it makes a water molecule become a hydronium ion H3 O+. Some acids dissociate better than others. HCl is a strong acid and it dissociates almost completely into H and Cl- ions. But CH3COOH (acetic acid - also called, vinegar) is a weak acid and doesn't dissociate as much into CH3COO- and H+ ions as does HCl. They both conduct electricity in water, but HCl is a much better conductor because it separates so well. +

The compounds which do dissociate into ions (soluble salts and acids), will give the water solution an ability to carry an electric current. You could put a light bulb into a solution of salt-water, connect it to a battery, and it would light. It would also work in an acid solution. But it would NOT work in a sugar solution. There are 3 big ideas in Solution chemistry that you need to know: 1. Understanding Molarity. Molarity is a fancy chemistry way of saying “concentration”. Molarity is Moles/Liters. It is the number of moles of a substance dissolved in liters of water. 2. Knowing how temperature and pressure affect how much solute you can dissolve in water. 3. Distinguishing electrolytes from non-electrolytes. Electrolytes are those compounds which dissociate into ions in solution. Non-electrolytes don’t. (see the overall summary of chapter above) Vocabulary you need to know: Substance: One element - for example: Gold, copper, etc. One compound - for example: Salt, sugar, CO2 Solution A solid, liquid or gas which has one substance evenly distributed in another. Solvent: The dissolving medium. (water is a typical solvent) Solute: The substance dissolved in the solvent. (in salt-water, salt is the solute) Homogeneous mixture: another name for a solution – usually when two solids are combined. Alloy: A homogeneous mixture (solution) of 2 or more elements with metallic properties (ex. Brass = Cu + Zn). Alloys are NOT chemically combined – there is no sharing of electrons. Miscible: Like dissolves like. Non polar substances will dissolve more easily in non polar liquids. And polar substances in polar liquids. Water is polar and will help dissolve polar compounds Heterogeneous mixture: When you can visibly see particles in solution. (sand/water) Solubility: The amount of a miscible substance that can be dissolved in another. See table C-10  on p. 920 for a listing of solubilities of ionic compounds in water. This

will let you know if a single or double displacement reaction will form a precipitate or not. If it is insoluble, that means it forms a precipitate. (Example: Sodium Chloride is soluble – NO precipitate) Precipitate: An insoluble substance in solution. Saturated/Unsaturated: A solution that cannot hold any more solute is saturated. Electrolytes: Substances which dissolve into water and allow it to carry an electric current (example: Salt water conducts electricity better than “pure” water). Non-Electrolytes: Covalent compounds are non-electrolytes. Sugar, for example, will not dissociate (it’s Covelent) into individual ions in water – therefore it cannot carry an electric charge A. Molarity Concentration: Measure of how much solute is dissolved in a specific amount of solvent/solution. ● A concentrated solution contains larger amount of solute ● A dilute solution contains a smaller amount of solute Molarity (M): moles of solute / liters of H2O ● Also known as molar concentration ● It’s unit is M (read/pronounced a molar) Example: 1) What is the molar concentration of a one liter solution with 0.5 mole of solute? Molarity = moles of solute / liter of solution = 0.5 mol solute / 1 L solution = 0.5 M 2) You are dehydrated and the doctor gives you an IV bag containing 100.5 mL solution containing 5.10g glucose (C6H12O6). What is the molarity of the solution? (Molar Mass of glucose = 180.16 g/mol) Molarity = moles of solute / liter of solution Step 1: convert grams of solute glucose to moles solute using Molar Mass of glucose 5.10 g glucose x 1 mole glucose = 0.0283 mole glucose 180.16 g glucose Step 2: convert mL of solution to Liters of solution (Remember week 1 notes: here we’re going from milliliters to liters so we’re going to a higher prefix this means we’re moving the decimal to the left 3 decimal places) 100.5 mL = 0.1005 L Step 3: solve for molarity M = moles of solute/ liters of solution = 0.0283 mole glucose / 0.1005 L solution = 0.282 M Homework #1: Solutions Worksheet page 1

B. Dilutions Dilutions: Adding additional solvent to dilute a more concentrated stock solution. ● The total # of moles of solute does not change during a dilution (only solvent is added) Dilution Equation: M1V1 = M2V2 M1 = the initial molarity of the concentrated solution V1 = the initial volume of the concentrated solution M2 = the final molarity of the dilute solution V2 = the final volume of the dilute solution Example: What volume of 2.00 M CaCl2 stock solution would you use to make 0.50 L of 0.300 M CaCl2? M1V1 = M2V2 M1 = 2.00 M

V1 = ?

M2 = 0.300 M

Solve the equation:

2.00M x V1 = 0.300M x 0.50L

V2 = 0.50 L Answer:

V1 = 0.075 L

Homework #2: Molarity and Dilution Problems worksheet Page 2 C. How Temperature and Pressure Effect Solubility p. 460 TEMPERATURE ➜As the temperature of a liquid increases, the solubility of a solid in that liquid increases EX. Hot tea will dissolve more sugar that cold tea. ➜As the temperature of a liquid increases, the solubility of a gas in that liquid decreases. EX. Warm soda loses its carbonation faster than cold soda. PRESSURE ➜As the pressure of a liquid increases, the solubility of both solids and gases in that liquid will increase. EX. When you open a can of pressurized soda, the pressure decreases and the bubbles escape. D. Net Ionic Equations p. 292-293 A net ionic equation lists the ionic formula for substances which are insoluble in water and just the individual ions for those substances which are soluble in water (see p. 920 for a listing of what ions are soluble in water and which are not). IMPORTANT – NOT ALL IONS ARE WATER SOLUBLE! Spectator ions: Ions which are soluble in water stay surrounded by water and thus are not part of the reaction and are called spectators. Spectator ions are dropped from the net ionic equation. See example below where spectator ions K and NO3 are dropped. Regular Equation: NaCl(aq) + AgNO3 (aq) + --> AgCl(s) + NaNO3(aq) Ionic Eq:

Na+ (aq) + Cl- (aq) + Ag+ (aq) + NO3- (aq) + --> AgCl(s) + Na+ (aq) + NO3 -(aq)

Net Ionic Eq: Ag+ (aq) + Cl- (aq) --> AgCl(s) Homework #3: pg. 294 #33-37 AND pg. 305 #90-94 (use page 920 for help with solubility)

E. Beer’s Law Recall what we learned about light earlier in the year. Colored solutions are absorbing some wavelengths of light. Beer’s Law tells us that as the concentration of a solution changes, so does the amount of light that is absorbed. Equation: A = ε M b A = Absorbance ε = absorptivity constant (otherwise known as the extinction coefficient) M = Molarity b = path length (this is generally 1) F. Molality m = molality = # moles of solute/kg of solvent Molality is used to predict temperature changes between boiling point and freezing points of liquids. DON’T GET THIS CONFUSED WITH OUR FRIEND, MOLARITY, M with is Moles/liter of H2O Practice Problem: How would you prepare a 0.50 m solution of NaCl in 500 g of water? 0.50 moles 1 kg

=

x moles 0.5 kg

x = 0.25 moles of NaCl convert moles to grams and add to 500 ml of water

* remember water’s density is 1 g/ml so 1000 ml is 1000 grams and 1000 grams = 1 kg G. Colligative Properties pg. 471 Colligative properties - A property of solutions that doesn’t depend on the size or type of molecule or atom in solution, just the concentration. For example, the Freezing point and boiling point of a liquid are determined by the number of particles in solution, not the type. 1 mole of sodium particles in 1L of water will do the same as 1 mole of sugar in 1L of water. This is very important. H. Calculating Freezing Point Depression and Boiling Point Elevation pg 472-475 Tb = kb x m x i

Tf = kf x m x i

T is the change of temperature from what it would normally be without any solute. When dealing with boiling, the Tb would be added to the original boiling point. When dealing with freezing, the Tf would be subtracted from the original freezing point. m is molality K is a constant used for the solvent you are talking about changing the boiling of freezing of. For example: kb for water is: 0.51 o C(kgH20))/mol kf for water is: 1.86 o C(kgH20))/mol  i is the number of ions that the solute dissolves into. For example, NaCl dissolves into two ions – Na+ and Cl- so i = 2 for NaCl. Aluminum nitrate (Al(NO3 )3) dissolves into one Al+3  and three NO3 - so i = 4. See Table 15-4 & 15-5 on pg. 472 & 474 for more k values of different solvents See Table C-12 on pg. 921

Fact: By adding a solute to any liquid, you increase   the boiling  point and reduce   the freezing  point. For example: adding salt to water will increase its boiling point and decrease its freezing point. Why does it increase the boiling point? You lower the vapor pressure of the liquid when you add a solute because the solute prevents the pure solvent molecules from escaping. (Recall that vapor pressure is the internal desire of a liquid to boil away). A lower vapor pressure means you have to heat water hotter to get it to vaporize. Salt water will boil at a temperature greater than 100o C. Why does it lower the freezing point? By adding a solute to a liquid solvent, you decrease the amount of intermolecular interactions between molecules of the pure solvent. Thus, if pure water can be frozen at 0o C, salt water will freeze at a colder temperature because the salt ions get in the way of the intermolecular attraction between water molecules. When salt water freezes, the frozen substance contains water only, not salt water – want proof? Icebergs floating in the ocean are pure water! More proof? How about this one: Spreading salt on snow melts it because the outside air temperature is cold enough to freeze pure water, but not cold enough to freeze salt water. What does this have to do with making ice cream? Ice cream is mostly frozen milk and sugar. You can think of ice cream as mostly water with some dissolved solute in there (lactose, protein…etc). This means that ice cream will freeze at a temperature BELOW 0o C (the normal freezing point of pure water.). Your freezer is kept at about –10o C. Cold enough to keep ice cream from melting. Pure ice (made from water) will draw heat out of the surrounding air to begin melting into liquid water. As  ice melts, its temperature remains constant until all of the ice has melted (we will learn more about this interesting phenomenon of solids keeping their temperature constant until they have completely melted in Chapter 16 – Thermodynamics) Normally, it will draw in enough heat to reach its melting/freezing point of 0o C. BUT…..If you add salt to the ice, this reduces the melting/freezing point of ice to about –10o C (remember adding a solute to a solvent will reduce the freezing point). Now, the ice will draw in enough heat to make its temperature –10o C. This is cold enough to freeze milk into ice cream. WOW, This is cool: Did you know that adding 1 mole of CaCl2 to water will drop the freezing point lower than if you added 1 mole of NaCl? Why is that? Because 1 mole of CaCl2 will break up into ONE mole of Ca+ cations and TWO moles of Cl- anions. But NaCl will break up into only ONE mole of Na+ ions and ONE mole of Cl- ions. The colligative property rule states that more ions makes a bigger change in the freezing point and boiling point. ONE mole of CaCl2 essentially becomes THREE moles when it breaks up. ONE mole of NaCl becomes TWO moles. How many moles does 1 mole of C12  H22O11 become when dropped into water? 1 mole. Why? Because it doesn’t break apart! It’s Covalent and it stays as one big molecule. Homework #4: Practice Problems # 33-36 p. 474

Pg 485 #76-79, 83, 86-88...