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Name: _____________________ Chemistry SOL Review Test Directions: Read each question and write a brief answer. Use your notes, diagrams, etc. to help you. SHOW ALL WORK (no work = no credit) Your work and writing must be neat & legible (unreadable = no credit) Unit 1: Measurement (Ch. 3) 1) Define the following accuracy reflects how close a measurement is to a) Accuracy a known or accepted value b) Precision c) Quantitative Data d) Qualitative Data precision reflects how reproducible measurements are, even if they are far from the accepted value quantitative analysis is used to tell 'how much' is in a sample data that determine the characteristics of a substance chemical constituents. 2) Number of Significant Figures a) 1.23 X 107 3 b) 100.1010 7 c) 33,000 2 d) 0.00004 1 3) Perform the calculations indicated. Round your answer to the correct number of significant figures. 1500 a) 632,000 410 b) 100.0 + 45.0 + 35 180. c) 98.012 +16 114 d) 891.000 * 1.0 890 4) Put the following numbers in proper scientific notation. -5 a) 0.0000330 3.30 x 10 b) 95,200,000 9.52 x 107 5) Take the following numbers out of scientific notation. a) 3.458 X 10-11 0.000 000 000 03458 10 b) 2.4 X 10 24,000,000,000 6) Answer the following: Be sure to put your answers in PROPER SCIENTIFIC NOTATION a) (3.10 x 103)+(1.68 x 102) 3.27 x 103 b) (1.11 x 108) – (1.1 x 103) 1.1 x 108 5 -9 c) (8.00 x 10 ) / (4.00 x 10 ) 2.00 x 1014 d) (600 x 1025)*(0.00020 x 10-10) 1 x 1014 7) Use the solving problems method to answer the following questions. SHOW ALL OF YOUR WORK. a) How many micrograms are in 3.00 grams? 3.00g 1,000,000 μg = 3.00 x 106 μg 1g b) How many seconds are in 24 days? 24 days 24 hrs 60 min 60 sec = 1 day 1 hr 1 min c) How many grams are in 9.42 mL of pentane (C5H12)? (Density = 0.0980 g/mL) 9.42 ml 0.0980 g = 1 mL d) How many milligrams are in 85 kilograms? 85 kg 1,000 g 1,000,000 μg = 1 kg 1g 2,073,600 sec 2,100,000 sec 0.923 g C5H12 8.5 x 1011 μg Unit 2: Matter & Energy (Ch. 2) 8) Classify the following as either a physical or chemical change a) Pizza is sliced Physical change b) Paper is torn Physical change c) NaCl is dissolved Physical change d) Wood burns Chemical change 9) ID the following as either a heterogeneous mixture or a homogeneous mixture. a) Italian dressing heterogeneous mixture b) NaHCO3 solution homogeneous mixture c) Muddy solution heterogeneous mixture 10) Predict whether the element is a metal, nonmetal, or metalloid. nonmetal a) selenium metal, b) magnesium nonmetal c) chlorine d) antimony metalloid 11) List four different characteristics of metals & nonmetals Nonmetal Metal Not Ductile (Brittle) Ductile Low MP and BP High MP and BP Tend to be volatile Luster (shiny) Insulators Conductors 12) Describe the following changes of state a) Sublimation Solid to Gas b) Freezing Liquid to solid c) Vaporization Liquid to Gas d) Condensation Gas to Liquid e) Melting Solid to Liquid 13) Differentiate between gases, liquids, & solids. solids liquids gases Slow Faster Fastest Particle speed close together no Far apart no Particle spacing Packed Volume Regular arrangement Fixed regular arrangement Fixed regular arrangement Variable (fixed or variable) Shape Definite Not Not (definite or not) Most dense Somewhat dense Least dense Density No Not easily Yes Compressible? 14) Identify the group/family number/s for the following: a) Noble gases 18 b) Halogens 17 c) The Alkali Metals 1 d) Alkaline Earth Metals 2 e) Transition Metals 3-12 Unit 3: Atomic Structure and Theory (Ch. 4) 15) Describe the Dalton’s atom model. Atoms are indivisible and indestructible. (Billiard ball model) 16) Describe Thomson’s atom model. Electron are like raisins surrounded by a soup of positive charge to balance the electrons' negative charges (Plum Pudding Model). 17) What experiment did Thomson do? The cathode rays tube 18) What was his model called? Plum Pudding Model 19) Describe Rutherford’s model. Atoms have a center nucleus and the rest of the atom is mostly empty space 20) What experiment did Rutherford do? He shot alpha particles through a thin film of gold and to his amazement a few alpha particles rebounded almost directly backwards. 21) What was his model called? “gold foil" experiment 22) Describe Bohr’s model. He proposed his quantized shell model of the atom, electrons orbits around the nucleus 23) What was his model called? Planetary Model 24) What is the difference between an isotope and an ion? Isotope Ion Different number of Neutrons Different number of Electrons 25) Use your periodic table to fill in the following table. (HINT: These are isotopes) Protons Neutrons Electrons Atomic Mass number number 33 44 33 33 77 a) As -77 153 74 54 53 127 b) I 30 35 30 30 65 c) Zn 2+ 26 30 24 26 56 d) Fe 5 6 5 5 11 e) boron-11 26) A cation is positive “+” in charge a anion is negative “‒“ in charge. 27) A horizontal row in the periodic table is called a Periods 28) A vertical column is called a Groups or Families H 29) proton B F G 30) isotope 31) electron 32) neutron D 33) nuclear pull E 34) shielding a. the smallest particle of an element that retains the properties of that element b. atoms of the same element with different masses c. the number of protons in the nucleus of an atom d. attraction of positive nucleus for the outer negative electrons e. electrons in between the outer electrons and the nucleus shield the positive nucleus from the negative outer electrons. f. the first subatomic particle discovered A C 35) Atom 36) atomic number g. subatomic particle with no charge h. .the kind of subatomic particle with a positive charge 37) The element iron is made up of the isotopes 58Fe (2.45%), 56Fe (80.3%), 54Fe (10.2%), and 57Fe (7.05%). Calculate the average atomic mass (weight) of iron. Answer: 58 (0.0245)+56(0.803)+54(0.102)+57(0.0705) = 55.92 Unit 4: Electrons in atoms (Ch. 5) 38) Describe the following prop 39) erties and their periodic trend (you may draw a picture for the trend) Definition Trend a) electron affinity → increases D H A C F E B G 40) 1st quantum # 41) p orbital 42) Atomic spectrum 43) Electromagnetic spectrum 44) 3rd quantum # 45) 2nd quantum # 46) d orbital 47) 4th quantum # a) the characteristic spectrum of an element b) complex shaped c) A listing of all known forms of radiation d) describes how far the electrons is from the nucleus e) describes the shape of the orbital f) describes the orientation of the orbital g) describes the direction the electron is spinning in h) dumbbell shaped 48) Which color has the shortest wavelength? ROY G BIV (longest to shortest so Violet) b) electronegativity → increases 49) Define the following terms. Frequency The amount of time it takes for one cycle to complete Amplitude The distance from the equilibrium to the crest or to the trough Wavelength The distance from Crest to Crest or Trough to Trough c) ionization energy → increases 50) Write the complete electron configuration for the following : 2 2 6 2 6 2 2 a) Titanium (Ti). 1s , 2s 2p , 3s 3p , 4s , 3d d) atomic radius ← increases e) ionic size (cations, anions) f) Atomic number g) valence electrons h) atomic mass i) electron configuration (s, p, d, f blocks) → increases → ↓ increases →increases → ↓ increases 2 2 6 2 6 b) Si-4 1s 2s 2p 3s 3p normally it would be 3p2 but there are 4 -4 additional electrons Si giving it a noble gas configuration Unit 5: Nomenclature (Ch. 9) 51) Name the following compounds: (molecular or ionic) HINT: Molecular = 2 nonmetals – use prefixes; Ionic = anything else – just state the name of the cation then anion (use roman numerals if it is a cation w/ multiple charges) a) CaS b) PO3 Calcium Sulfide Phosphorus trioxide c) N2O6 d) (NH4)3CO3 Dinitrogen hexoxide Ammonium carbonate e) C2I4 f) Fe2O3 Dicarbon tetraiodide Iron (III) oxide 52) Write the formulas for the following compounds b) rubidium carbonate a) lead (IV) chloride Rb2CO3 PbCl4 c) silicon monoxide d) sodium arsenide SiO Na3As e) strontium nitrate Sr(NO3)2 Unit 6: The Mole (Ch. 10) f) silver nitride Ag3N Use the solving problems method to answer the following questions. SHOW ALL OF YOUR WORK and use significant figures. 53) Find the molar mass of Sr3(PO3)2 Sr 87.62(3)+ P 30.97(2) + O 16.00(6) = Answer: _420.80_ 54) How many formula units are in 2.84 moles of potassium oxide (K2O)? 2.84 mol K2O 6.02 x 1023 fu = Answer: 1.71 x 1024 fu K2O 1 mol K2O 55) How many moles are in 8.50 X 1023 molecules of water? 8.50 x 1023 mlcs H2O 1 mol H2O = Answer: 1.41 mol H2O 6.02 x 1023 mlcs H2O 56) How many grams are in 4.42 moles of methane (CH4)? 4.42 mol CH4 16.05 g CH4 = Answer: 67.4 g CH4 1 mol CH4 57) How many moles are in 1.2 X 104 liters O2 gas at STP? 1.2 x 104 L O2 1 mol O2 22.4 L O2 = Answer: 540 mol O2 58) How many grams are in 0.9500 liters F2 (Fluorine Gas)? 1 mol F2 38.00 g F2 = Answer: 1.612 g F2 0.9500 L F2 22.4 L F2 1 mol F2 59) How many atoms of nitrogen are in 56.652 grams of nitrogen? 56.652 g N2 1 mol N2 6.02 x 1023 mlc N2= Answer: 1.2171 x1023 mlc N2 28.02 g N2 1 mol N2 60) Find the percentage composition of a compound that contains 19.5 g of iron and 5.8 g or sulfur. The total mass of the compound is 20.3 g. Answer: omit_ 61) 11.66 g iron, 5.01 g oxygen. Find the empirical formula. 11.66 g Fe 1 mol Fe 55.85 g Fe 5.01 g O 1 mol O 16.00 g O = 0.20877 / 0.20877 = 1 x 2 = 2 = 3.13125 / 0.20877 = 1.5 x 2 =3 So the empirical formula is Fe2O3 62) What is the molecular formula? empirical formula NO2, molar mass = 92.02 g/mol N 14.01(1) +16.00(2) = 46.01 92.02/46.01 = 2 So there are 2 (NO2); so the molecular formula is N2O4 63) What is used to convert molecules to moles? a) Name: Avogadro’s number b) Number: 6.02 x 1023 mlcs/mole Unit 7: Chemical Equations (Ch. 11) 64) Use the activity series of metals to determine which of the following reactions will occur. If a reaction will take place, complete and balance the equation. If the reaction will not occur, write no reaction. a) 2 Al + 3 CuSO4 3 Cu + Al2(SO4)3 b) 6 Ag + 2 H3PO4 3 H2 + 2 Ag3PO4 65) Determine the products of the following double replacement reactions. a) 3 Ag2CO3 + 2 Na3PO4 2 Ag3PO4 + 3 Na2CO3 b) Pb(NO3)2 + CoCl2 PbCl2 + Co(NO3)2 66) What are the meanings of the following symbols? a) (aq) aqueous b) ↑ given off (gas) c) ↓ Precipitate (solid) d) ∆ heat 67) Directions: Balance the following reactions. To the right of each equation, write the type of reaction each represents. (Use S, D, S-R, D-R, C rather than the name) a) Zn + HCl → ZnCl2 + H2 Single Repl b) C7H6O2 + O2 → CO2 + H2O c) Sb + O2 → Sb4O6 Combustion Synthessis d) H2O2 → H2O + O2 Decompostion e) Al2(SO4)3 + Ca(OH)2 → Al(OH)3 + CaSO4 Double Repla Unit 8: Stoichiometry (Ch. 12) 68) Define the following: It is the reactant that is completely used up a) Limiting Reagent b) Actual Yield c) Percent Yield d) Theoretical Yield in a reaction The experimental yield, what you actually get. (Actual yield/Theoretical yield) x 100. Percent yield measures how efficient the reaction is under certain conditions. The amount of product that could possibly be produced in a given reaction. Answer: _________ 73) Sodium chloride can be prepared by the reaction of sodium metal with chlorine gas. 2Na (s) + Cl2 (g) 2 NaCl (s) a) When 6.70 mol of Cl2 reacts with 3.20 mole of Na, what is the limiting reactant? Answer: __________ b) How many moles of NaCl are produced? Problems: Make sure that your chemical equation is balanced. Set up a proportion under the substances in the equation that you are interested in. Divide by the coefficient to cancel out moles, the (coefficient * molar mass) to cancel out grams, or the (coefficient * 22.4) to cancel out liters 69) How many moles of aluminum oxide are needed to react completely with 5 moles of SnO? 2Al + 3SnO 3Sn + Al2O3 5 mol SnO 1 mol Al2O3 = 1.7 mol Al2O3 3mol SnO 70) How many moles of sodium are needed to react with 18.3 g of Sodium chloride? 2 Na + Cl2 2 NaCl 18.3g NaCl 1 mol NaCl 2 mol Na = 0.313 mol Na 58.44 g NaCl 2mol NaCl Answer: _________ c) How much of the excess reagent remains unreacted? Answer: _________ 74) Calcium carbonate is decomposed by heating. CaCO3 CaO + CO2 What is the percent yield of this reaction if 27.8 g of CaCO3 is heated to actually produce 13.6 g of CaO? Answer: _________ 71) If 48.3 g of Fe2O3 react with Aluminum in a single displacement reaction, how many grams of aluminum are needed? 2 Al + Fe2O3 Al2O3 + 2 Fe 48.3g Fe2O3 1 mol Fe2O3 2 mol Al = 0.605 mol Al 159.7 g Fe2O3 1mol Fe2O3 75) 2 Al + 2 H3PO4 2AlPO4 + 3H2 What is the percent yield if 6.6 g of Al is reacted to produce 1.85 g of H2? 72) If 40 L of oxygen combine with oxygen to synthesize water, how many grams of water are produced? 2H2 + O2 2 H2O Unit 9: GASES (Ch. 14 &Ch 13) (don’t forget to change C to K) 76) What do moles tell you? What is the unit for n? Answer: _________ 77) What is volume? What are the units of volume? 78) What are the units for pressure and how are they related? 1 atm = 101.325 kPa = 760 mmHg = 760 torrs 79) Fill in the following table GayAvogadro’s Charles’ Boyle’s Lussac’s Formula V1 = V2 P1V1=P2V2 P1 = P2 V1 = V2 T1 T2 n1 n2 T1 T2 Inverse or Direct Directly Directly Directly Inversely Graph R= 8.31 L*kPa mol*K ; 0.0821 L*atm mol*K ; 62.4 L*mmHg mol*K 80) Many gases are available for use in the laboratory in compressed gas cylinders, in which they are stored at high pressures. Calculate the number of moles of O2 that could be stored at 24.0 C and 0.935 atm, in a cylinder with a volume of 16.5 liters? PV=nRT Answer: _______________ 81) A balloon contains 2.5 mol of helium at a pressure of 1.06 atm and a temperature of 274. K. What is the volume of the balloon? PV=nRT Anwer: _______________ 82) A helium balloon with a volume of 22.5 L, a pressure of 101.3.kPa, and a temperature of 328 K is put into an environment where the pressure is 25.0 kPa and the temperature is 25.0C. What is the new volume of the balloon? Answer: _______________ 83) The gas balloon has a volume of 4.8 L at 330 K. At what temperature will the balloon expand to 8.50 L? Answer: _______________ 84) A 2.43 L balloon contains 0.35 moles of He. How much Helium must be added in order for the balloon to be 5.2 L? Assume constant temperature and pressure. Answer: _______________ 85) Calculate the temperature (C) of 2.65 g of I2 if it occupies 50. mL at 676 mm Hg. PV=nRT 676 mmHg 0.050 L = (2.56g/253.8 g/mol) (62.364 L mmHg mol-1 K-1) T 53 K or -219 C Answer: _______________ Unit 10: Liquids (Ch. 15&16) 86) The solubility of most liquids increases as temperature increases. (increases/decreases) 87) The solubility of gases in liquids increases as temperature decreases. 88) Three colligative properties are: 1. Vapor pressure depression 2. Boiling point elevation 3. Melting point depression 4. Osmotic pressure 89) Four factors which affect solubility. a) b) c) d) Temperature Pressure Solvent – Solute interaction (like and unlike) Size 90) What is the molarity of a 1500 mL solution made with 7.5 moles of KCl? 7.5 mol KCl = 1.6 M 1.5 L 91) How many mL of a 0.01 M solution of NaCl would you need in order to have 1.4 moles? 1.4 mol = 0.01 M NaCl x = 140,000 mL xL 92) How could you make a 2.5 M NaOH dilute solution in a 250 mL volumetric flask, from a 3 M concentrated solution of NaOH. M1V1 = M2 V2 2.5 M 250 ml = 3 M x x = 208 mL So place 208 mL of the 3M solution in a flask and add 42 mL of H2O for 250 mL of 2.5 M solution. 102) Fill in the following chart Lewis Dot Diagram CO # of lone # of Molecular Bond VSEPR pairs bonds Geometry Angles Structure (name) skip skip H2Se 93) If NaCl dissolves in water? Is NaCl POLAR or nonpolar? Unit 11: Bonding (Ch. 8 & 7.3) 94) Write electron dot structures for the following ATOMS a) silicon b) S 95) Write electron dot structures for the IONS of the following atoms a) S-2 GeH4 AlF3 b) Mg+2 96) Use the electronegativity table to determine what type of bond (IONIC, VERY POLAR COVALENT, MODERATELY POLAR COVALENT, OR NONPOLAR COVALENT) will form between the atoms of the following elements. IONIC 97) Mg and F MODERATELY POLAR COVALENT 98) C & H POLAR COVALENT 99) C and O IONIC 100) Na & F MODERATELY POLAR COVALENT 101) H and Br Unit 12: RATES OF REACTION & CHEMICAL EQUILIBRIUM (Ch.18) 103) Using the following equations for complex ions, draw an arrow to show how each system would change under the following conditions. (In which direction will the equilibrium shift?) [Zn(H2O)4]2+ + H2O ↔ [Zn(H2O)3(OH)]+ + H3O+ a. Removing water _________ b. Adding H3O+ _________ 104) For the following reaction, equilibrium is established at a certain temperature when the following concentrations are present: [CO] = 0.10 mol/L, [H2O] = 0.80 mol/L, [CO2] = 0.012 mol/L, and [H2] = 0.012 mol/L. Calculate the Keq value for this reaction. HINT: Write the equilibrium expression for the reactions, then fill in the values and solve for Keq CO (g) + H2O (g) ↔ CO2 (g) + H2 (g) Answer: ________________ Unit13: Acids, Bases, & Neutralization (Ch. 14) 105) In a neutralization reaction the products are always _______ and _______. 106) An acid turns litmus paper ______ and a base turns litmus paper _______. 107) Water is amphoteric because it can _____________________ 108) pH = ___________[H+] 109) Label the Bronsted-Lowry acid, base, conjugate acid, and conjugate base in the following reactions. a) H2O + H3PO4 H2PO4-1 + H3O+ + b) H2O + H2SO4 H3O + HSO4 -1 110) Name each of the following acids and their anions Formula Name of anion Name of acid a) HBr b) H2CO3 c) H3PO3 d) H2SO3 111) Write the formulas for the following acids. Name Cation Anion Formula a) Chloric acid b) Phosphorous acid c) hydronitric acid 112) Complete and balance the following neutralization reactions. a) H2SO4 + Ca(OH)2 ______ + ________ b) HI + NH4OH _______ + ________ 113) Determine the pH for the aqueous solutions with the following [H+]. Is the solution acidic, basic or neutral? [H+] pH Acid, base, or neutral a) 1 X 10-6 M b) 1 X 10-7 M 114) Determine the [H+] for aqueous solutions that have the following pH values. Is the solution acidic, basic, or neutral? pH [H+] Acid, base, or neutral a) 11 b) 8 115) Determine the pH of the following solutions that have the following [OH-]. Is the solution acidic, basic, or neutral? [OH-] pH Acid, base, or neutral a) 3.8 X 10-3 b) 0.000001 116) Would a stronger acid have a large or a small Ka? 117) In a titration, a 15 mL solution of hydrochloric acid is neutralized by 6 mL of 2.5 M sodium hydroxide, using phenophthalein as an indicator. What is the concentration of the hydrochloric acid solution? Answer: _____________ Unit 14: REDOX (Ch. 20) O.I.L. R.I.G. 118) Oxidation is the loss / gain of electrons & increase/decrease of oxidation #. 119) Reduction is the loss / gain of electrons & increase/decrease of oxidation #. 120) Show the oxidation numbers for all of the atoms in the following equation. Al(OH)3 + H2SO4 ---> Al2(SO4)3 + HOH 121) In the following equation, which chemical is oxidized? Reduced? Oxidizing agent? Reducing agent? Zn + HCl ---> ZnCl2 + H2 Oxidizing Agent: __________ Reducing Agent: __________ Oxidized: _________ Reduced: _________ 129) Write the half reactions for the oxidized and reduced chemicals in previous problem. Oxidation ½ reaction : ____________________________ Reduction ½ reaction: ____________________________ Unit 15: Thermochemistry (Phase Changes)(Ch. 17) Heat energy = mcT (WITHIN A PHASE) Heat Energy = m H (TO CHANGE PHASE) 4.18 J = 1 Cal & specific heat of liquid water = 1 cal/gC 130) A 16.3 gram water sample in a calorimeter has its temperature raised 8.0 C while an exothermic chemical reaction is taking place. How much heat is generated in the calorimeter? (solve for calories and for joules) Answer: ____________ Answer: ____________ 131) How many joules would be needed to raise 20.0 grams of water forms 10.0C to 98C? Answer: ____________ 132) If 5.0 grams of water at 45C gains 5600.0 Joules of heat, what is the final temperature of the water? (Hint - solve for T) Answer: ____________ 133) 98.7. grams of a material is melted at it’s freezing point. Heat of fusion = 472 J/g Answer: ____________ 134) 16.3 grams of a material is heated from 100oC to 200oC Specific heat = 2.34 J/goC How much heat is absorbed? ___ 122) the energy required to melt of mole of a substance ___ 123) the energy required to change one mole of a substance from a liquid to a gas. ___ 124) a measure of the randomness or disorder of a system ___ 125) The total energy content of a system ___ 126) A substance that increases the rate of a chemical reaction and that is not changed by the reaction. ___ 127) The system gains energy ___ 128) The system releases energy a. entropy b. heat of vaporization c. heat of fusion d. endothermic e. exothermic f. enthalpy g. catalyst Answer: ____________ Match the following: 135) Decide which of these reactions are endothermic and exothermic,Your DONE!! _____ a. BCl3 + 3 H2O H3BO3 + 3 HCl H = -112kJ _____ b. 2 Fe + 3 CO2 + 26.8 kJ Fe2O3 + 3CO _____ c. C2H4 2 H2 + 52.3kJ _____ d. S2Cl2 + CCl4 CS2 + 3 Cl2 H = +112kJ