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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 2. ____ Na2CO3 (aq) 4. ____ helium 6. ____ H2O (l) 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. __________ burning gas 3. __________ decompose water into hydrogen and oxygen 4. __________ evaporate water Units of measurement cm3, L or mL is __________________ g is _______________________ g/cm3, g/l or g/mL are ________________ joules , calories, kcal is ______________________ K or ºC 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) 18.63 g + 5.2 g 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? Conversions Given: 1 joule = 0.239 calorie. Convert 630 joules to calories. 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. Symbol Atomic Number Protons Xe Neutrons Electrons Mass Number 77 Ba-137 27 33 27 Na+ 55 22 Mn 25 17 15 18 36 O-2 Write the complete chemical symbol for the ion with 31 protons, 39 neutrons, and 28 electrons. 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 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 3. 78.3 g Pb to atoms Final Exam Review Day 2 Name_________________ Electron Configuration (shows where all the ______________ are located) Diagonal rule 1s2 2s2 2p6 3s2 3p6 3d10 2 6 4s 4p 4d10 4f14 5s2 5p6 5d10 5f14 2 6 6s 6p 6d10 6f14 7s2 7p6 7d10 7f14 Draw the electron configuration for the following: Potassium Sulfur Cobalt P−3 Mg+2 Valence electrons are electrons in the outer shell. Draw the dot diagrams for the following: Potassium P−3 Sulfur Mg+2 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 3-13 Group 18 ___________________ ___________________ ___________________ Group 2 Group 17 _____________________ _____________________ The Noble gases are stable because they have ___ ______________ ____________. Elements are placed in order of ______________ _______________ 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 __________. 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 a) b) c) d) e) f) g) h) i) j) k) l) m) n) o) F K C Ba Si N Mn Mg Ge Fe Be Ca Na Rb Au Se Na O B I Cs Cl S Br Sb F Mg I Sr Pb 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 Chemical Bonding The three types of chemical bonds are 1. _________________ usually between a ___________ and a ___________. There is a __________ of electrons. 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 ______________. 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 or polyatomic ion. Predict the shape names. Predict the intermolecular force for the molecules. CCl4 SiO2 BaCl2 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 Name_________________ 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 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 % composition by mass of each element 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 82 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) 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 F2 NO H3PO4 (aq) NaClO3 SO2 Recognize an acid, a base a salt and other: SO2−2 MgSO4 _____ HCl _____ C6H12O6 _____ 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 → + O2 (g) → Classify and balance the following. ________ 1. BaCl2 + (NH4)2CO3 → KCl ________ 2. KClO3 ________ 3. Na2O + P4O10 ________ 4. C6H6 + O2 ________ 5. FeCl2 (aq) + + → → → BaCO3 + NH4Cl O2 Na3PO4 CO2 Na3PO4 (aq) + H2O → Fe3(PO4)2 (s) + NaCl (aq) 3 CO2(g) + 4 H2O(g) What do the symbols (s), (l), (g) and (aq) mean? Which substances in question 5 are solutions? What are the reactants in question 5? What are the products in question 5? What is the subscript of oxygen in question #4? What is the coefficient of oxygen in question #4? Stoichiometry 1. Given the following: C3H8(g) + 5 O2(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? (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? Final Exam Review Day 4 Name_________________ Gas Laws Kinetic Molecular Theory assumes gases are made up of _________ ___________ moving in _____________ ___________, colliding into each other with ______________ collisions. Ideal vs Real : Non-ideal behavior occurs 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 1 atm (760 mm Hg) and 273 K 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 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. Calculate the number of grams of N2 in a 6.0 liter cylinder at 27C and 800 torr. Gas Stoichiometry 1. Given: 2 H2(g) + O2(g) → 2 H2O(g) If this reaction occurs at 200C 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 500C and 850 torr are produced from 6.60 g C3H8? Phases and Phase Diagrams Solids Definite shape Definite Volume Fluidity Particle movement (KE) Particle arrangement Attractive forces Density Diffusion rate Compressibility Liquids Gases 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°C to 30°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 Final Exam Review Day 5 Concentration % by mass = g solute × 100 g solution Name_________________ Molarity (M) = moles solutes Liter solution 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 g Mg(OH)2 is dissolved in 80 g H2O. 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? 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? 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: 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 2. Calculate the pH of a 0.0052 M solution of HCl. pH = − log (N of acid) 3. 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. 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: N1V1 = N2V2 How many mL of 0.15 N KOH are needed to neutralize 50.0 mL N H2S, if both completely break into ions? Final Exam Review Day 6 Name____________________ 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? Heat in reactions Refer to the following diagram for the hypothetical reactions: E n e r g y E n e r 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) (b) What is the value of ∆H for the reaction? 3. (a) (b) Oxidation and Reduction Oxidation = _______ of electrons Reduction = _______ of electrons The oxidizing agent caused the oxidation. (It was reduced) The reducing agent caused the reduction. (It was oxidized) KNO3 + CO → CO2 + NO2 + K2O Assign oxidation numbers to the elements in each substance. ___________ was oxidized. ___________ was reduced. ___________ was the reducing agent. ___________ was the oxidizing agent.