CHEM 150 Section 01, Q2 2016 Midterm 1 Student name... Student number... Time: 90 Mins ATTEMPT ALL THE QUESTIONS 1 Formulae and constants pv = nrt P 1 V 1 T 1 = P 2V 2 T 2 Ptotal = p1 + p2 + p3 +... U = q + w H o = Σ ν p ΔH f o (products) Σν r ΔH f o (reactants) At standard temperature and pressure (1 atmosphere = 101.325 kpa, T=273.15K) 1 mole of an ideal gas occupies 22.4 L R = 8.3145 L kpa mol -1 K -1
1. State whether each of the following is (i) a pure material; (ii) a homogenous mixture or (iii) a heterogeneous mixture: Black coffee or undissolved sugar in it) Homogeneous mixture (provided that you haven t got any grains Sand Heterogeneous mixture, since it contains a variety of different substances, as well as bits of shell, tar,... compounds Gasoline Homogeneous mixture; it s a solution of many different Carbon nanotubes Pure material 2. Categorise each of the following elements as (i) metal, (ii) non-metal or (iii) metalloid: Aluminium metal Argon non-metal Element number 16 (sulfur) non-metal metalloid, on the boundary between metals and nonmetals Germanium 2 3. In a reaction between phosphorus and hydrogen, 6.319 g of phosphorus required 0.308 g of hydrogen for complete reaction. Determine the molecular formula of the compound formed, showing your working. To find the formula we need to determine how many moles of each substance react. We can use the formula # moles=weight/aw The atomic weight of hydrogen is 1.0, so the MW of H2 is 2, and that of P is 30.97. Using the formula we find that the number of moles of P is 6.319/30.97 = 0.204, and of hydrogen is 0.308/2.0 = 0.154. The ratio of P to H2 is then approximately 4:3, or P to H (atoms) is 4:6, so the formula could be P4H6, or P2H3 (we cannot distinguish between these without further information). 4. What is an isotope? Isotopes are atoms of the same element that contain different numbers of neutrons (and are consequently of different mass). Since they are both the same element, isotopes have the same number of protons in the nucleus, and have the same number of electrons; they have virtually identical chemical properties.
5. Beside each of the following statements, write down whether the statement is true or false. (i) The ionisation energy of Mg is greater than that of Mg +. False. We d expect Mg + to hold onto its electrons more strongly than the neutral atom because of the positive charge. Ionisation energy always increases with each electron we remove from an atom. (ii) The electron affinity of chlorine is less than the electron affinity of sodium. False. Chlorine is keen to acquire an extra electron so it can gain a noble gas configuration. By contrast, sodium wants to lose an electron to reach the same state. (iii) Oxygen atoms are larger than phosphorus atoms. False. Size increases as we move down a group, and also as we move from right to left along a row. On both counts we d expect oxygen to be smaller than phosphorus. (iv) In the HCl molecule the electrons in the bond are pulled towards hydrogen, as it is the more electronegative element. False. Chlorine is the more electronegative element, so it will have a greater share of the electrons in the covalent bond. (v) Neither Na + nor Mg 2+ have any valence electrons. True. Sodium, being in group I, started with one valence electron before it was removed, and magnesium, in group II, started with two valence electrons. 3 6. The electron affinity of chlorine is -349 kj mol -1. The ionisation energy of sodium is 495 kj mol -1. (a) Define electron affinity and ionisation energy. Electron affinity is the amount of energy released when an electron is added to an atom to form a negative ion, all species being in the gas phase. Ionisation energy is the energy required to remove an electron from an atom to form a positive ion, all species being in the gas phase. (b) Calculate the energy required to ionise a single sodium atom (note that Avogadro s number is 6.022 x 10 23 ). 495000/6.022 x 10 23 = 8.22 x 10-19 J (c) Calculate the energy released when one electron is added to a single chlorine atom. Similarly, - 5.80 x 10-19 J
(d) Sodium chloride forms stable crystals containing ionic bonds. Including in your answer a comment about the values found in parts (b) and (c), explain why sodium chloride is stable. Energy is released when an electron joins Cl, but is required to remove an electron from Na. The former is a smaller value than the latter (see the answers to parts (b) and (c)), so overall energy is required to transfer one electron from sodium to chlorine. The reason why NaCl is stable despite this is that there is strong electrostatic attraction between the two ions (especially since, in the solid, each positive ion is surrounded by six negative ions and vice versa) and this formation of ionic bonds is more than sufficient to make the overall process of formation of the solid favourable. 7. The molecules shown above differ by only a single atom, and have similar molecular weights and shapes. Which of the two molecules should have the higher boiling point? Justify your answer. 4 The N-species can form intermolecular Hydrogen-bonds (i.e., between one molecule and another) because nitrogen is very electronegative. Phosphorus is less electronegative than nitrogen, so does not form hydrogen bonds. When a liquid boils, intermolecular bonds must be broken. The stronger these bonds are the higher the temperature that is required for boiling, so, since the stronger intermolecular bonds are in the nitrogen-containing molecule, this should have the higher boiling point. 8. For each of the following types of intermolecular force, (i) explain the origin of the force, (ii) give an example of a molecule in which the force is important: (a) Dispersion force Dispersion forces arise when the semi-random movement of electrons in atoms or molecules creates a temporary dipole (because the centre of the negative charge due to the electrons no longer coincides with the centre of the positive charge due to the nuclei). This temporary dipole in one atom/molecule will create a similarly short-lived dipole in a neighbouring molecule.
These dipoles are oriented in such a way that they always attract, so dispersion forces always tend to bind molecules to each other, albeit weakly. The forces exist in every atom and molecule, but are most important in species which otherwise would not interact much, for example argon atoms. (b) Hydrogen bonding When hydrogen is covalently bonded to a very electronegative atom (specifically F, O or N), the electrons in the bond are shared very unevenly, with the hydrogen being significantly positive. This creates a bond dipole. The hydrogen in such a molecule is then attracted to the F, O or N in a neighbouring molecule (or, sometimes, to F, O or N in another part of the same molecule) and, since the partial charges on the two atoms are quite large, a weak bond forms. The strength of such a bond is typically only around 10 kj per mole, compared to a typical covalent bond energy of a couple of hundred kj, but the H-bonds are sufficiently strong to cause easily measurable effects in the molecule, such as an increase in the boiling point. Examples include HF, H2O, NH3 and CH3COOH. (c) Dipole-dipole interactions If a molecule has a permanent dipole (HCl, or CH3Cl for example) or has no molecular dipole, but does have bond dipoles (CO2 for example) the dipoles in two different molecules can attract each other, again giving rise to weak intermolecular bonds. 5 9. What is meant by the terms open, closed and isolated systems in thermochemistry? Both energy and matter can enter or leave an open system. In a closed system only energy can be transferred between system and surroundings, while in an isolated system neither energy nor matter can enter or leave. 10. The following standard molar enthalpies of combustion have been measured (note that the formula of acetylene is C2H2,): Acetylene Carbon (graphite) Hydrogen -1299.6 kj -393.5 kj -285.8 kj (a) Define standard molar enthalpy of combustion.
The enthalpy change when one mole of material is formed from the stable form of the elements, all materials in this standard state. (b) Write down and, if necessary, balance the three chemical equations to which the figures given above apply. 2C2H2 + 5O2 4CO2 + 2H2O (1) H: 2 x (-1299.6) C + O2 CO2 (2) H: -393.5 2H2 + O2 2H2O (3) H: 2 x (-285.8) (c) Calculate the standard enthalpy of formation of acetylene. This is the enthalpy change for the reaction 2C (s) + H2(g) C2H2(g) With all species in their standard states. We can make this equation from the equations above: 2 x (2) + ½ x (3) ½ x (1), so the enthalpy change is 2 x (-393.5) + (-285.8) (-1299.6) = +226.8 kj 6 11. Draw a circle around each of the thermochemical functions listed below which are state functions: Entropy Work Free energy Enthalpy Temperature Heat Those in red are state functions. 12. (a) In thermochemistry what is a spontaneous process? A process which can continue without outside intervention. (b) What happens to the entropy (of the universe) during a spontaneous process? It goes up. The entropy of some part of the universe, such as the system we are concentrating on, may fall, but this must be compensated for by a rise in entropy elsewhere. (c) Which thermodynamic law is expressed by your answer to part (b)?
The 2 nd law of thermodynamics. (d) The change in Gibb s Free Energy, G, during a chemical reaction is given by G = H T S. Comment on whether the processes shown below would be spontaneous (a) at low temperature, (b) at high temperature, (c) at all temperatures, or (d) never. Justify your answers. 2NH3 3H2 + N2 for which H and S are both positive. The H and T S terms have opposite signs, so the sign of the free energy change depends on temperature. At low temperature H will dominate and G will be positive, so the reaction will be unfavourable. At high temperature the entropy term will dominate and the process will be favourable. 3 O2 2 O3 for which H is positive and S is negative. Both terms, H and T S, are positive, so the reaction is always unfavourable. 13. The vapour pressures of toluene (C7H8) and benzene (C6H6) at 25 o C are 28.4 torr and 95.1 torr respectively. A solution is prepared that contains 50g of benzene and 12g of toluene. (a) What is meant by the term torr? 7 A Torr is a unit of pressure, equal to the pressure exerted by a column of mercury 1 mm high. One atmosphere is, by definition, equal to 760 Torr. (i) 92 (ii) 78. (b) Calculate the molecular weight of (i) toluene and (ii) benzene. (c) Determine the mole fraction of (i) toluene and (ii) benzene in the solution. Moles benzene = 50/78 = 0.64. Moles toluene = 12/92 = 0.13. Total moles 0.64 + 0.13 = 0.77. xb = 0.64/0.77=0.83; xt=0.13/0.77=0.17. Check xb+xt=1.0 ok. (d) Calculate the partial pressures of (i) benzene and (ii) toluene; hence calculate the total pressure above the solution. pa = pa o. xa pbenz = 95.1 x 0.83 = 78.9 Torr. PTol = 28.4 x 0.17 = 4.8 Torr. Total 78.9 + 4.8 = 83.7 Torr.
14. Hexane and pentane (right) are mutually soluble in all proportions. Water is virtually insoluble in both liquids; explain both observations. Liquids may be mutually soluble provided that the types of intermolecular bonding are not too different. The only significant intermolecular bonds in hexane and pentane are the weak bonds caused by dispersion forces. Hexane molecules can easily displace pentane molecules, and vice versa, since the energy of the new intermolecular forces in the mixture is comparable to the energy of the forces in the separate liquids. By contrast, if hexane were to dissolve in water, the hexane molecules would have to break the strong H- bonds in water, but would then be able to interact with water molecules only with weak dispersion forces. It is therefore energetically unfavourable for hexane to dissolve in water. 8 15. Water gas, which is an equimolar mixture of hydrogen and carbon monoxide, can be made by reacting water vapour with hot carbon: H2O (g) + C (s) H2 (g) + CO (g) H= +131 kj mol -1 (a) Is the reaction written above endothermic or exothermic? Endothermic, since the enthalpy change is positive (i.e., heat is being absorbed by the system). (b) In a typical experiment 70 kg of carbon react with water vapour. How many moles of product are formed? Check: is the equation balanced? Yes. 1 mole of carbon creates one mole of each of hydrogen and carbon monoxide. Moles carbon=weight/aw = 70000/12 = 5833, so there are 5833 moles of each product, or 11666 moles in total of product.
(c) Use the Ideal Gas Law to calculate the total volume of the product, assuming that the pressure is 200 kpa and that the temperature is 740K. pv=nrt, so V=nRT/p = 11666 x 8.314 x 740/200 = 358867 litres (d) Assuming (i) that the temperature is constant during the reaction, and (ii) that the volume of carbon can be neglected, calculate the change in volume of gas during the reaction. Easy. The equation shows that 1 volume of gaseous reactant (water) makes two moles of product, so the change in volume is half of the final volume, i.e. 179434 litres. Since this is a drop in volume, we should give it a minus sign: -179,434 litres. 16. During a process a system absorbs 23kJ of heat from the surroundings and does 78 kj of work on the surroundings. Which of the following statements are true and which are false? The energy of the system increases by 101 kj. False The surroundings lose 55 kj of energy. False 55 KJ of energy are transferred from system to surroundings. True The energy of the system falls by 55 kj. True It is not possible to know whether the energy of the system changes. False 9