Chemistry 30S Unit 1 States of matter Physical Phase-States of Matter 1. Solid: high density, hard to expand/compress and rigid in shape. 2. Liquid: mid-high density, hard to expand/compress but takes shape of container. 3. Gas: low density, ease to expand / compress to fill container. 4. Plasma: a special gaseous state where particles are highly ionized (charged). Takes the shape of its container but is also influenced by magnetic fields. Plasma is the most common state in the universe, but uncommon on Earth (except as lightning). A phase is a region of uniform properties. Changes in State: Phase Change Associated Term Phase Change Associated Term Solid Liquid Melting Liquid Solid Freezing /Crystalization Liquid Gas Boiling Gas Liquid Condensing Gas Solid Deposition Solid Gas Sublimation *Note: DISSOLVING is a special type of change involving a state of matter (solid, liquid or gas) mixing with a liquid substance such that their particles are evenly spread in-between each other. If a substance dissolves in water the solution formed is said to be AQUEOUS. **Note: Endothermic process for melting, boiling, and sublimation Exothermic process for freezing, condensing, deposition.** eating Curves and Cooling Curves: Both heating curves and cooling curves show the following: 1. Melting temperature (melting point)/ Freezing temperature (freezing point) 2. Boiling temperature (boiling point) / Condensation temperature (condensation point) 3. State of matter (solid, liquid, gas) 4. Energy changes (in form of heat/temperature)
eating Graph of Ethanol Temperature (celsius) 140 120 100 80 60 40 20 0-20 -40-60 -80-100 -120-140 -160-180 Liquid 0 5 10 15 20 25 30 1 st plateau = melting/freezing temp Solid Time Gas phase 2 nd plateau = boiling/condensation temp *Note: At the plateau, a phase change occurs. The substance is in a mixed or ETEROGENEOUS PASE. The heat energy is used to break and rearrange bonds between atoms. Example: at the first plateau, the substance is in solid and liquid states. Once all the transition between states is complete, the temperature of the substance will increase once again. There is a dynamic equilibrium between the two phases. **Note: At the slopes, the substance is being warmed up. Its molecules are vibrating, but no phase change occurs here thus, the substance is in a single, or OMOGENOUS PASE. ***Note: For a cooling curve, the general shape of the graph is the same as a heating curve, but the graph starts at a higher temperature and works its way down to a lower temperature. By using heat arrows, explain the following phase changes. 1. Liquid to solid Endothermic or Exothermic 2. Solid to gas Endothermic or Exothermic 3. Gas to liquid Endothermic or exothermic 4. Solid to liquid Endothermic or exothermic 1.Explain why melting is an endothermic process. 2. Explain with diagrams how ice can boil water. b. Why does it have to be in a closed system? 3. What does it mean to reach a plateau during a heating/cooling curve?
4. What is an elastic collision? 5. ow does the kinetic energy of particles vary as a function of temperature? 6. Explain why the baking instructions on a box of cake mix are different for high and low elevations. Would you expect to have a longer or shorter cooking time at a high elevation? 7. Compare and contrast vaporization and evaporation 8. Explain the relationships amount vapor pressure, atmospheric pressure and boiling point. 9. Use the kinetic-molecular theory to explain why gases are easier to compress than liquids or solids. 10. What is a dynamic equilibrium? Give an example. 11. Be able to explain various phenomena such as boiling water with ice, evaporative cooling, and cooking in higher altitudes. 12. Material Boiling ( C) - becomes a gas Freezing ( C) - becomes a solid 2 O (water) 100 C 0 C Fe (iron) 2750 C 1535 C O (oxygen) -183 C -218 C g (mercury) 357 C -39 C Find out the state for each of the temperatures Material -499C -183C -5C 53C 504C Fe Ethyl Alcohol 78 C -114 C O g Ethyl Alcohol Water 13. Explain why ionic solids have a higher melting point than covalent solids. 14. Describe the difference between the densities of the states of matter. Include examples to support your answer. 15. What does kinetic mean in terms of kinetic energy and the movement of molecules?
16. Explain the relationship between intermolecular forces and vapour pressure. Unit 02: Gas Laws GAS LAWS are a series of formulas/equations that describe the effects of volume, pressure and temperature on gases. About Gases: 1. Gases are considered the simplest state of matter 2. Gases are composed of molecules in random order 3. Molecules collide elastically 4. Molecules collide with walls of container to exert pressure 5. Increasing temp, increases KE applied to the molecules giving them energy to move away from each other. Questions: 1) What is the atmosphere and what elements make up the atmosphere? 2) Convert the following: a) Convert 0.875 atm to mmg b) Convert 745.0 mmg to atm c) Convert 0.955 atm to kpa d) Convert 98.35 kpa to atm 3) A gas occupies 1.56 L at 1.00 atm. What will be the volume of this gas if the pressure becomes 3.00 atm? 4) 500.0 ml of a gas is collected at 745.0 mm g. What will the volume be at standard pressure? 5) 300 ml of O 2 are collected at a pressure of 645 mm of mercury. What volume will this gas have at one atmosphere pressure? 6) 600.0 ml of air is at 20.0 C. What is the volume at 60.0 C? 7) A gas occupies 1.00 L at standard temperature. What is the volume at 333.0 C? 14) At 210.0 C a gas has a volume of 8.00 L. What is the volume of this gas at -23.0 C? 15) A gas syringe contains 56.05 milliliters of a gas at 315.1 K. Determine the volume that the gas will occupy if the temperature is increased to 380.5 K 16) If 15.0 liters of neon at 25.0 C is allowed to expand to 45.0 liters, what must the new temperature be to maintain constant pressure? 17) Determine the pressure change when a constant volume of gas at 1.00 atm is heated from 20.0 C to 30.0 C. 18) A 30.0 L sample of nitrogen inside a rigid, metal container at 20.0 C is placed inside an oven whose temperature is 50.0 C. The pressure inside the container at 20.0 C was at 3.00 atm. What is the pressure of the nitrogen after its temperature is increased? 19) A sample of gas at 3.00 x 10 3 mm g inside a steel tank is cooled from 500.0 C to 0.00 C. What is the final pressure of the gas in the steel tank?
20) 93.0 ml of O 2 gas is collected over water at 0.930 atm and 10.0 C. What would be the volume of this dry gas at standard conditions? 21) 690.0 ml of oxygen are collected over water at 26.0 C and a total pressure of 725.0 mm of mercury. What is the volume of dry oxygen at 52.0 C and 800.0 mm pressure? 22) If a gas is heated from 298.0 K to 398.0 K and the pressure is increased from 2.230 x 10 3 mm g to 4.560 x 10 3 mm g what final volume would result if the volume is allowed to change from an initial volume of 60.0 liters? 23) A gas balloon has a volume of 106.0 liters when the temperature is 45.0 C and the pressure is 740.0 mm of mercury. What will its volume be at 20.0 C and 780.0 mm of mercury pressure? 24) What is SCUBA and why is it important to slowly acclimatize while resurfacing to the water? 25) Conceptual questions: Explain the ammonia fountain and how it works Explain how you can boil water with a vacuum Explain what happens when you boil water in different altitudes Explain how a aluminum can be crushed from inverting into a bowl of cold water Explain the Cartesian diver mechanism Unit 3 Chemical reactions 1) Name the following: a) N 2 O 5 b) CuBr 2 c) PbSO 4 d) Na 2 CO 3 e) CCl 4 f) K 3 PO 4 2) Balance the following equations: a) Al + O 2 Al 2 O 3 b) Potassium chlorate when heated produces potassium chloride plus oxygen gas 3) Identify the types of reactions in the following: a) CO 2 C + O 2 b) NaCl + AgNO 3 NaNO 3 + AgCl c) Zn + CuSO 4 ZnSO 4 + Cu d) 2O 2 + N 2 N 2 O 4 4) Name all the diatomic molecules. 5) Predict the products of the following: a) AgNO 3 + NaCl b) Zn + Cl 6) Find the formula mass for the compounds listed in question 1. 7) ow many molecules are found in 0.750 mol of zinc? 8) Calculate the number of moles in 16.9g of carbon tetrachloride
9) Calculate the mass of 3.80 moles of sodium carbonate 10) Determine the number of moles in 33.6L of e gas at STP 11) Given Zn + 2Cl ZnCl 2 + 2 find: a) ow many moles of hydrogen are produced with 3 moles of zinc? b) ow many litres of zinc chloride are made given 4L of hydrochloric acid at STP? c) ow many grams of hydrochloric acid are present with 56g of zinc? d) ow many litres of hydrogen are produced at STP given 30g of hydrochloric acid. e) ow many molecules of zinc are present with 2 mol of zinc chloride? 12) If we had 2.5 mol of 4 and 4.8 mol of O 2, according to the following reaction, find: 4 + 2O 2 CO 2 + 2 2 O a) Which reactant is limiting? b) ow many moles of CO 2 would be produced? c) ow much of the excess is left over? 13) If we had 2.0 mol Cl reacting with 2 moles of MgO, find the following using the reaction: 2Cl + MgO MgCl 2 + 2 O a) Which is the limiting reagent? b) ow much of the excess is left over? c) ow many grams of MgCl 2 would be made? 14) Boron has two naturally occurring isotopes with masses of 10.0129 amu which occupies 19.91 percent and another isotope of 11.0093 amu and occupying 80.09 percent. Calculate the average atomic mass of Boron 15) Determine the empirical formula of methane given that 6.0 g of methane can be decomposed into 4.5 g of carbon and 1.5 g of hydrogen. 16) The composition of a compound is 40% sulfur and 60% oxygen by weight. What is its empirical formula? Copper reacts with sulphur to form copper (I) sulfide according to the following balanced equation: 2 Cu (s) + S (s) Cu 2 S (s) What is the limiting reagent when 40.0 g Cu reacts with 12.5g S? Unit 4 - Solution concepts & theory ow are pure substances different than mixtures? Think in the atomic level. Are compounds such as NaCl classified as pure substance or mixtures? Give three other types of compounds and 3 other types of mixtures (any type) Give 2 examples of a heterogeneous and homogenous mixture. ow are solutes different than solvents? When can a mixture not be a solution? What is it meant by aqueous solution? Give an example If you have two liquids dissolve in each other, which one is considered to be the solute and which one is solvent? What are the 9 types of solutions and give an example of each. What are the 4 properties of a solution? ow does temperature affect the solubility of a solid dissolved in water? ow does pressure affect the solubility of a gas dissolved in water? What is freezing point depression and how do we achieve that?
Polarity What is boiling point elevation and how do we achieve that? What are ways of expressing concentration? What is polarity and how does it play in solutions? Give an example of a polar and non-polar molecule. Water is polar, list two other examples that can be mixed in water. ow can you determine if a compound is polar, non polar or ionic? Saturated, unsaturated, and supersaturated solution. What two processes can you do to saturate a solution? What about supersaturating a solution? ow can you UNSATURATE a SUPERsaturated solution? What happens when you cool down a supersaturated solution and why? Reading solubility curves What does the line represent? What does the area above and below represent? On the solubility curve, what is the solubility expressed as? If you were assigned the task of finding the solubility curve of MgCl2, how would you do it? What are you measuring.? Know the calculations from the worksheets. Write dissociation equations for the following electrolytes dissolving and dissociating in water. Show the physical state of all species involved. a) potassium hydroxide b) sodium carbonate c) potassium permanganate d) ammonium sulfate Questions to calculate 1. Suppose you made a saturated solution by adding 67.5g of solute into 200mL of water. What is the solubility of this substance? Ans: 2. Suppose you made a saturated solution by adding 450g of solute into 0.5L of water. What is the solubility of this substance? Ans: 3. The solubility of substance B is 0.92g/mL. a) If you had 125mL of water, how much substance B could you dissolve? Ans: b) If you had 7 kg of substance B, what is the minimum amount of water needed to dissolve all the solute (to make a saturated solution)? Ans:
c) If 40.0g of substance B was added to 20.0mL of water, what type of solution would exist: unsaturated, saturated or supersaturated? Molarity Ans: 4. Using a diagram, draw how this ionic compound can be dissolved by water: NaO (Solvation process) 1. A solution of sodium carbonate, Na 2 CO 3, contains 133g of solute in 623mL of solution. a) What is the concentration? b) What is the molarity? 2. Boris has an NaCl solution with a concentration of 8g/L. a) If Boris poured 0.175L into a separate beaker, what is the concentration? b) What would be the molarity (mol/l) of the NaCl in the beaker? 3. A mass of 178g of sulfuric acid ( 2 SO 4 ) is dissolved in water to make a 0.50M solution. What volume of water is needed to make this solution? 4. ow many grams of glucose (C 6 12 O 6 ) is dissolved in a 500mL solution, if the molarity is 1.30M? 5. If a solution of 2 SO 4 was said to be 7.5% w/v, if you had 500mL of this acid, what is the mass of 2 SO 4? ow many moles is that? 6. Becky has a 86%v/v solution of vinegar. The bottle is about 600mL. ow much vinegar is actually Becky s bottle? 7. Describe how you would prepare 6.00 L of 0.500 M K 2 CrO 4 from 2.75 M K 2 CrO 4 stock solution. 8. Describe how you would prepare 50.0ml of a 1.0M NaO solution with all the available items in the laboratory. 9. A solution is prepared by dissolving 3.8 g potassium chlorate in enough water to make 100.0 ml stock solution. A 10.00-mL sample of this stock solution is then placed in a 50.00-mL volumetric flask and diluted to the mark with water. What is the molarity of the new solution? 10. A stock solution of Cl is 6M, you need to create a diluted bottle (1L) of 0.1M of Cl. Calculate and describe the steps. 11. ow would you prepare 25 ml of 0.010M NaCl starting with a 0.500 M NaCl solution? 12. If I have 440 ml of a 1.5 M NaBr solution, what will the concentration be if I add 560 ml more water to it? 13. If I dilute 150 ml of 0.50 M lithium acetate solution to a volume of 750 ml, what will the concentration of this solution be? Solubility Curve Questions: Use the solubility curve (attached) to help you answer the following questions: 1. Which substance has the greatest solubility?
2. Which substance(s) tend not to increase its solubility with increasing temperature? 3. Which substance is the least soluble at 80 C? 10 C? 4. At what temperature will potassium chlorate and sulfur dioxide both have the same solubility? 5. What is the solubility of ammonia (N 3 ) 10 C? 90 C? 6. At what temperature will potassium nitrate and ammonia both have the same solubility? 7. Is there a temperature where 3 substances have the same solubility? ow can you tell? 8. Suppose you made a saturated solution of Cl at 45 C. ow many grams of Cl was needed to make this solution? 9. Suppose you have a saturated solution of KClO 3 at 100 C. ow many grams of solute would precipitate if you somehow removed 50mL of water? 75mL? 100mL? 10. At 70 C, you poured 160g of potassium nitrate into beaker containing 100mL of water. What kind of solution do you have? 11. In the previous question, how much precipitate would be seen at the bottom of the beaker? ow much precipitate would there be if the temperature dropped to 15 C? 12. If the temperature was set at 50 C, rank all the substances from greatest solubility to lowest solubility. 1. Name the following organic molecules using the proper IUPAC system of naming: Structural formula IUPAC Name 1.
2. C 3 C 3. 2 C 4. 5. 2 C 6. 3 C
2. Name the following hydrocarbons and state what class of hydrocarbon does each belong to (alkane, alkene, alkyne, aromatic, or cyclic). Structural Formula Class IUPAC Name Cl Cl Br Br Additional Review Questions from All Sections Please answer on a separate sheet! (2009) 1. Magnesium has three isotopes, Mg-24, Mg-25, and Mg-26. The percentage of each in order is 78.70%, 10.13%, 11.17%. Calculate the relative (apparent) atomic mass of magnesium. 2. Using the Kinetic Molecular Theory, explain what happens to the molecules of a solid as it changes phase to a gas as temperature is increased? 3. In order to turn 1kg of ice at -30 C into steam at 110 C, what is the total amount of heat (in kilojoules) that needs to be absorbed? Cp of ice = 2.06J/gC; Cp of water = 4.18J/g C; Cp of steam = 2.02J/gC; fus = 334J/g; vap = 2260J/g (Based on this info, include a sketch of the heating curve.)
4. There are 9.76 x 10 23 particles of Br 2 gas in a container. a) ow many moles are there? b) What volume would this gas occupy? 5. At STP, you are given the following gases: 2 moles Br 2 and 2 moles I 2. Which gas would occupy the greater volume? Which gas would have the greater mass? 6. In this reaction: 3 2 (g) + N 2 (g) 2N 3 (g) What volume of N 3 measured at STP is produced when 2.15 L 2 reacts? 7. Calculate the empirical formula of a substance with 46.55% Fe and 53.45% S. 8. You are given the following percent compositions: C = 40.0%, = 6.7%, O = 53.3%. a) What is the empirical formula? b) If the molecular mass of this substance is 180g/mol, what is the molecular formula? 9. In the this reaction: Al 2 (SO 4 ) 3 + 6 NaO 2 Al(O) 3 + 3 Na 2 SO 4 ; If 285g of NaO and 256.5g of Al 2 (SO 4 ) 3 are reacted in a container: a) Which substance is the limiting reagent? b) ow many grams of aluminum hydroxide will be produced as a result? c) ow moles of sodium sulfate will be produced as a result? d) ow many grams of the excess reactant will be left as a result? 10. Suppose you have a 0.05M stock solution of Cl. If you poured out 150mL and diluted it with 50mL of water, what would be the new molarity?