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Acl Chem Final Exam Review - Day 1 Name __________________ Classify the following matter as an element, a compound or a mixture. 1. ____HCl 3. ____ Cl2 (g) 5. ____ barium 2 ____ HCl (aq) 4. ____ diet coke 6. ____ calcium nitrate Classify the following matter as a pure substance or a mixture 1. ____ Na2CO3 (s) 3. ____ air 5. ____ dirt 7. _____ MgCl2 2. ____ Na2CO3 (aq) 4. ____ helium 6. ____ H2O (l) 8. _____ tea Classify the following matter as homogeneous or heterogeneous. 1. ____ tossed salad 3. ____ stainless steel 5. ____ Ca(OH)2 (aq) 2. ____ Cl2 4. ____ saline solution 6. ____ MgSO4 (s) Classify the following as a physical change or a chemical change. 1. __________ dissolving sugar into water 2. __________ decompose water into hydrogen and oxygen 3. __________ burning gas 4. __________ evaporate water Classify the following as matter or energy 1. ____ air 2. ____ light 3. ____ heat 7. _____ Ne Classify the following as a metal, nonmetal or metalloid. 1. ____ Na 2. ____ Kr 3. ____ As 7. _____ H 6. ____ brittle 7. ____ ductile 8. ____ insulator 9. _____ has luster Units of measurement cm3, L or mL is _____________________ g is _______________________ g/cm3, g/l or g/mL are ________________ K or ºC is __________________ joules , calories, kcal is _______________ Temperature conversions ºC = K – 273 20º C = __________K Metric Conversions ________ mm = 1 m ___________ cm = 1m __________ m = 1 km Significant Figures 1. Reading Instruments Properly – give all figures known and estimate the last one 2. Determine the number of significant figures in the following measurements. 3. a) 87 000 000 000 __________ b) 0.000 607 0 __________ c) 320.00 __________ Calculate each of the following. Report answers using significant figures. a) (4.15 × 105) m × (6.024 × 1023) m b) __________3.56_g__________ 3.6 cm × 2.5 cm × 5.2215 cm c) 18.60 g + 5.20 cg Density D = M /V Calculate the density of an object measuring 21 cm by 6.0 cm by 2.12 cm and having a mass of 522.2 grams A piece of granite has a mass of 55 g. When I place it into a graduated cylinder that has 50.0 mL of water in it the level changes to 70.0 mL. What is the density of this piece of granite? % Error % Error = (Experimental – Actual) ÷ Actual × 100 The density of a sample of water was experimentally determined to be 1.08 g/ml. What is the % error? 2 Conversions Given: 1 joule = 0.239 calorie. Convert 630 joules to kilocalories. Given: 2.24 lb = 1 kg. Convert 18 lbs to kg. Subatomic Particles and their Jobs # protons = atomic number (identifies the element) neutrons define the mass of the atom # protons + neutrons = mass number # electrons determines the charge Fill in the following table: Subatomic particle Relative Mass (amu) Charge Location in atom 0 0 1 Fill in the following table: Symbol Atomic Number Protons Xe Neutrons Electrons Mass Number 77 Ba-137 27 33 27 Na+ 55 Mn 15 22 25 O-2 Write the complete chemical symbol for the ion with 31 protons, 39 neutrons, and 28 electrons. How many neutrons are in chromium−52? 3 Iron consists of four natural isotopes: Isotope Mass (amu) Percent Abundance 54 Fe 53.9696 5.82 56 Fe 55.9349 91.66 57 Fe 56.9354 2.19 58 Fe 57.9333 0.33 Calculate the average atomic mass of iron to five significant digits. Convert the following 1. 45.6 g Mg to atoms 2. 45.9 moles gold to g Final Exam Review Day 2 Electron Configuration (shows where all the ______________ are located) Diagonal rule 1s2 2s2 2p6 2 3s 3p6 3d10 4s2 4p6 4d10 4f14 2 6 5s 5p 5d10 5f14 6s2 6p6 6d10 6f14 2 6 7s 7p 7d10 7f14 Draw the electron configuration for the following: Potassium Sulfur Cobalt P−3 Mg+2 Valence electrons are electrons in the outer shell. Determine the number of valence electrons and draw the dot diagrams for the following: Potassium Sulfur P−3 Mg+2 4 Draw the shortcut electron configuration for each of the following Potassium Sulfur Cobalt P−3 Mg+2 Identify the following [Xe]4f 46s2 contains 3 electrons in its sixth and outer main energy level the element that has 2 electrons in the p sublevel in its second main energy level 4s24p5 Be able to locate s, p, d, and f blocks on the periodic table The ___________ ____________ _____________ is the same as the period number. There are ________ main energy levels. The ____________ ___________ _________ is caused when electrons emit energy as they fall back to a lower energy level. Periodic Table Know the name and location of the following groups Group 1 ___________________ Group 2 _____________________ Group 3-12 ___________________ Group 17 _____________________ Group 18 ___________________ The Noble gases are stable because they have ___ ______________ ____________. Elements are placed in order of increasing ______________ _______________ and placed in groups according to their ___________________ __________________. Rows = ______________ =___________ __________ _______ Columns =______________ = _______________ Metals are located _________________________________ Nonmetals are located______________________________ Metalloids are located __________________________________ Metals _____________ electrons. They form __________________. Nonmetals ______________________ electrons. They form __________. 5 Trends Atomic radius (size) ionization energy (energy to remove an electron) Electronegativity (ability to gain an electron) Reactivity for metal (metallic character) Reactivity for nonmetal Ionic Radius a) b) c) d) e) f) g) h) i) j) k) l) m) n) o) p) F K C Ba Si N Mn Mg Ge Fe Be Ca Na Rb Au Ca+2 Se Na O B I Cs Cl S Br Sb F Mg I Sr Pb K+1 is more reactive. is the smaller atom. has the higher ionization energy. has only s electrons in its outer shell has 4 valence electrons is a metal last electron is d5 goes to a charge of −2 has higher electronegativity is a metalloid is a halogen is more reactive is more likely to combine with oxygen loses electrons more readily is a transition metal has large ionic radius Chemical Bonding The three types of chemical bonds are 1. _________________ usually between a ___________ and a ___________. There is a __________ of electrons. The bond is due to attraction between _______________ and _______________. 2. _________________ usually between a ___________ and a ___________. The electrons are _____________. a) _____________ _______________ between 2 same nonmetals. Electrons are ____________ shared. b) _____________ _______________ between 2 different nonmetals. Electrons are ____________ shared. 3. _________________ usually between a ___________ and a ___________. Many electrons are shared. They are ______________. Classify the following chemical bonds as ionic, covalent or metallic: _____ CuO _____AlBr3 _____ H2O _____ PCl5 _____ MgSO4 6 The type of bond is determined according to the difference in ___________________ between the two atoms. |----|---------------|-------------------| 0 .3 1.7 4 Characteristics of ionic vs covalent. Ionic bonds are strong. Ionic substances have orderly pattern and a high ____________ ___________. They form ______________. Lewis Structures Predict the type of bonding. Draw Lewis structures for molecules. Predict the shape of the molecule. CCl4 SiO2 O2 H2O CO3−2 NH3 What is an intermolecular force? What are the three types of intermolecular forces? What is the strongest type? Final Exam Review Day 3 Calculate the oxidation numbers of the atoms in each of the following compounds or polyatomic ions: H3AsO4 H2Cr2O7 TiO2 SO3−2 Calculate the molar mass of the following 7 CaCl2 (NH4)2CO3 Mn3(PO4)2 Convert each of the following. a. 5.6 × 1026 molecules NaOH to grams NaOH b. 45g KClO3 to moles KClO3 c. 7.4 moles Sn(SO4)2 to molecules d. 3.25 moles (NH4)2CO3 to grams Find the % of phosphorus in Fe3(PO4)2. A compound contains 68.04 g N and 155.52 g O. Find its empirical formula. A compound contains 68.04 g N and 155.52 g O. Its molar mass is 92 g. Find its molecular formula. Naming Compounds and writing formulas Binary Ionic MgCl2 CuCl2 Tertiary Ionic MgSO4 Covalent CO2 Binary acids HI (aq) Oxyacids HNO3 (aq) HNO2 (aq) 8 Write the formula for each of the following substances: Strontium hypobromite Phosphorus trichloride Sodium chromite aluminum sulfite Magnesium nitrite chloride Calcium sulfide iron (II) hydroxide Sulfuric acid hydrosulfuric acid Phosphorus trioxide hydrogen Name the following substances HCl (aq) CaO SO2 F2 NO SO2-2 H3PO4 (aq) NaClO3 Recognize an acid, a base a salt and other: MgSO4 _____ C6H12O6 _____ HCl _____ NaOH _____ Chemical Reactions: Synthesis (element + element → compound) Mg (s) Decomposition (compound → element + element) HgO (s) → Single replacement (element + cpd→element + cpd) K + Ni(NO3)2 → Double replacement (cpd + cpd → cpd + cpd) Pb(NO3)2 + NaBr → Combustion (CxHy + O2 → CO2 + H2O) C3H8 + O2 → Neutralization (acid + base → salt + water) LiOH + H2SO4 → Classify and balance the following. ________ 1. BaCl2 + (NH4)2CO3 → KCl ________ 2. KClO3 ________ 3. Na2O + P4O10 ________ 4. C6H6 + O2 ________ 5. FeCl2 (aq) + → + → → + O2 (g) → NH4Cl O2 Na3PO4 CO2 Na3PO4 (aq) BaCO3 + + → H2O Fe3(PO4)2 (s) + NaCl (aq) 9 Identify the precipitate in question #5. What are the products in question 5? What do the symbols (s), (l), (g) and (aq) mean? What is the subscript of oxygen in question #4? Which substances in question 5 are solutions? What is the coefficient of oxygen in question #4? What are the reactants in question 5? Predict whether the following single replacement reactions will occur. 1. ________ Ca 2. ________ NaCl 3. ________ Li + AgNO3 I2 → K2O → + + → Predict if the following double replacement reactions will occur. 1. ________ Na2CO3(aq) + Hg(NO3)2(aq) → 2 NaNO3(aq) + HgCO3(s) 2. ________ Ba(NO3)2(aq) + MgCl2(aq) → BaCl2(aq) + Mg(NO3)2(aq) 3. ________ Ca(OH)2(aq) 4. ________ FeS(s) + 2 HCl(aq) → H2S(g) + FeCl2(aq) + H2SO4(aq) → Stoichiometry 1. Given the following: C3H8(g) + 5 O2(g) → 3 CO2(g) + 4 H2O(g) a) How many moles of C3H8 are needed to produce 9 moles of CO2? b) How many grams of CO2 are produced when 2.0 moles of C3H8 react with excess oxygen? c) How many grams of oxygen are needed to react with 22.0 g C3H8? d) How many grams of oxygen are needed to react with 3.48 liters of C3H8 at STP? 2. 3. Given the following: 2 Na + Cl2 → 2 NaCl (a) If 23.0 g Na and 71.0 g Cl2 react, what is the theoretical mass of NaCl produced? Why? What is the limiting reactant? What is the excess reactant? b) An excess of sodium is reacted with 142 g Cl2. If 199 g of NaCl are actually produced, what is the percent yield? Aluminum reacts with hydrochloric acid according to the following unbalanced equation: 2 Al + 6 HCl → 2 AlCl3 + 3 H2 (a) If 18 g Al are combined with 75 g HCl, which reactant is the limiting reactant? Which is the excess reactant? (b) What mass of AlCl3 is theoretically formed? (c) An excess of HCl is reacted with 52 g of Al. If 5.0 g H2 are actually produced, what is the percent yield? 11 Final Exam Review Day 4 Gas Laws Kinetic Molecular Theory assumes gases are made up of _________ ___________ moving in _____________ ___________, colliding into each other with ______________ collisions. As temperature increases, the particle movement also _____________________. Gases do not behave ideally when gases stop moving (or move slowly) That is at _________ temperatures and __________ pressures. P1V1 T1 = P2V2 T2 PV = nRT Ptot = P1 + P2 + P3 … STP means ______________ (760 mm Hg) and _______________ 1. A sample of gas occupies 400.0 mL at a pressure of 1 atm. What will the volume be if the pressure is changed to 2.00 atm while the temperature remains constant? 2. Calculate the volume occupied by 64.0 grams of O2 at a pressure of 850 torr and a temperature of 25C. 3. A gas occupies 250 mL at STP. It expands to 375 mL as it is heated to 75ºC. What is the new pressure? 4. A gas sample is collected over water when at a pressure of 760 mm Hg. What is the pressure of the dry gas if the partial pressure of the water vapor is 11.5 mm Hg? 5. A 600. mL sample of gas at STP expands to 900. mL at constant pressure. What is the new temperature? 6. Calculate the number of grams of N2 in a 6.0 liter cylinder at 27C and 800. torr. 12 Gas Stoichiometry 1. Given: 2 H2(g) + O2(g) → 2 H2O(g) If this reaction occurs at 200.C and 1500 torr, how many liters of O2 will react with 20.0 L of H2? 2. Given: C3H8(g) + 5 O2(g) → 3 CO2(g) + 4 H2O(g) How many liters of CO2 at 500.C and 850 torr are produced from 6.60 g C3H8? Phases and Phase Diagrams Solids Liquids Gases Definite shape Definite Volume Fluidity Particle movement (KE) Particle arrangement Attractive forces Density Diffusion rate Compressibility 13 Heat Calculations Q = m∙Cp∙∆T Q = m∙∆H 1. What quantity of heat, in joules, is required to raise the temperature of 22.5 g of lead from 18.0°C to 30.0°C? The specific heat of lead is 0.129 J/g·°C. 2. Calculate the quantity of heat energy (to the nearest calorie) required to convert 50.0 g of ice at 0°C to steam at 110°C. Be sure to draw a diagram showing the temperature and phase changes and the heat energy increases. Specific heats: water = 1.00 cal/g·°C steam = 0.480 cal/g·°C ∆Hvap = 540 cal/g; ∆Hfus = 80 cal/g 14 Final Exam Review Day 5 Concentration % by mass = g solute × 100 g solution Molarity (M) = moles solute Liter solution Molality(m) = moles solute Kg solvent 1. Calculate the molarity of 95.0 grams of Al(OH)3 dissolved in 500. mL of solution. 2. Calculate the % of a solution containing 35 g NaCl in 110 g H2O. 3. Calculate the % concentration by mass of a solution in which 20.0 g Mg(OH)2 is dissolved in 80.0 g H2O. 4. What is the molality of a solution made by dissolving 2.5 g Ca(OH)2 in 250 g of water? 5. How many g of NaCl are needed to make 2000. mL of a 2.50 M solution? 6. Find the freezing point depression and the boiling point elevation of each of the following ELECTROLYTES in solution. a. 18.0 g MgCl2 in 250. g water b. 80.0 NH4 Br in 100. g water 15 Solubility 1. What is the solubility of KCLO3 at 30 ºC? 2. 50 g of KCl are added to 100 g of water at 40°C. Is the solution saturated or unsaturated? If saturated, how many more grams of KCl remain undissolved? If unsaturated, how many more grams of KCl could be dissolved? 3. How does temperature affect the solubility of a solid in a liquid? 4. How does temperature affect the solubility of a gas in a liquid? 5. How does pressure affect the solubility of a solid in a liquid? 6. How does pressure affect the solubility of a gas in a liquid? Dilutions: C1V1 = C2V2 How many mL of a 12.0 M HCl solution are needed to make 100 mL of a 1.0 M HCl solution? 16 Ions in solution 1. Write an equation showing the dissociation of AlCl3 in water. 2. For the following write (a) the ionic equation; and (b) the net ionic equation. formula equation: 2 AlCl3(aq) + 3 Na2CO3(aq) → Al2(CO3)3(s) + 6 NaCl (aq) ionic equation: net ionic equation: What is the precipitate in the equation above? 3. Write an equation showing the ionization of H3PO4 in water. 4. Which of the following are electrolytes? CaBr2 CH4 Ca(OH)2 Na3PO4 HF Acids and Bases Be able to identify a neutralization reaction. 1. Calculate the Normality of: Normality (N) = M × Subscript of H+ or OH− a. 0.45 M H2SO4 b. 0.82 M Ca(OH)2 17 2. pH = − log (N of acid) Calculate the pH of : a. 0.0052 M HCl 3. b. 0.001 N HNO3 In an aqueous solution if H3O+ = 1 10 −5, then OH− = ________, pH= ______, and pOH =______. Know pH scale pH + pOH = 14 |-----------------|-----------------| acid ←7 → base Acids donate a proton. Bases accept a proton. Show a single step of the ionization of the following acids in water. Label the conjugate acid-base pairs. (a) HCrO2-1 + H2O → (b) HBr + H2O → Indicate each of the following as an acid, base, or salt: H2SO4 KBr HNO3 Ba(OH)2 Na2SO4 Properties of acid 1. Properties of base 1. 2. 2. 3. 3. 4. 4. 5. 5. Titration: NAVA = NBVB How many mL of 0.15 N KOH are needed to neutralize 50.0 mL 0.25 N H2S, if both completely break into ions? 18 Final Exam Review Day 6 Equilibrium Keq= Products reactants 1. A(g) Given: + 2 B(g) → AB2(g) + heat Write the equation for the equilibrium constant, Keq. If at equilibrium the concentrations are: [A] = 1.2 × 10−2 M; [B] = 2.0 × 10−2 M; [AB2] = 1.5 × 10−5 M ; calculate the value of the equilibrium constant. Is the reaction endothermic or exothermic? 2. 3. For the same equation given above, predict the effect of the following (get more products, get more reactants, or no effect on the amount of products or reactants) Adding more B Removing some A Heating decreasing the pressure Adding a catalyst removing AB2 Ksp = products The higher the Ksp the greater the solubility. For the reaction: Al(OH)3 (s) → Al+3 (aq) + 3 OH−1(aq) Write the expression for the solubility product, Ksp. 4. Name 4 ways to increase reaction rate. 5. What must you have in order for a reaction to occur? 19 Heat in reactions Refer to the following diagram for the hypothetical reactions: 70 E n e (kJ) 40 r 20 g y 70 E 60 n e (kJ) r 20 g y Time 1. Time Is the reaction endothermic or exothermic? (a) 2. (b) What is the value for the activation energy for the forward reaction? (a) 3. (b) What is the value of ∆H for the reaction? (a) (b) Classify the following processes as endothermic or exothermic. 1. __________ clothes dry 2. __________ condense gas 3. __________ a bomb explodes 4. __________ H2O(g) + ____________∆H = -314.43 kJ 5. C(s) + 131.3 kJ → CO(g) + H2(g) Oxidation Numbers Assign oxidation numbers to the elements in each substance H2SO3 HNO3 P2O5 MnO4−2 20