Is sodium hydride an Arrhenius base
WS 2010/11 (Prof. U. Schilde)
Update after the respective lecture
Lecture 1 (21.10.2010)
1. What is chemistry about?
2. How can you classify fabrics?
3. Assign the following substances according to the classification:
H2O, S8, Saline, air, amalgam, SO2, Milk, blood, caustic soda, hemoglobin!
4. Are the following processes chemical reactions or physical?
Operations? Give reason!
a) Dissolving sugar
b) decomposition of water
c) melting glass
d) rusting iron
f) boiling water
g) baking cakes
h) distillation of petroleum
i) Milk souring
j) hardening (setting) of mortar
5. Name six basic physical operations with which heterogeneous systems become homogeneous
and let solutions separate into pure substances and explain them briefly with an example!
Lecture 2 (22.10.2010)
Important parameters (pdf file for download) Stoichiometry - instructions for self-study (pdf file for download)
6. What is meant by the amount of substance, the AVOGADRO number, the molar mass, the molar volume
as well as under the molar concentration and the equivalent concentration?
7. Complete the following table!
Molar mass / g • mol-1
Mass of this amount of substance / g
8. 250 g of an aqueous potassium bromide solution with a mass fraction of 12% are to be prepared.
How much is to be weighed? [30 g KBr and 220 g H2O)
9. A bottle with acetic acid is labeled: 85% (volume fraction of pure acetic acid).
How many ml of pure acetic acid are in 800 ml of this solution? [680 ml]
10. What is the molar concentration of a sodium hydroxide solution of the 800 ml 3.2 g sodium hydroxide
contain? [0.1 mol / l]
11. What mass of zinc sulfate (ZnSO4 x 7 H2O) is used to produce 0.8 l of a 0.1 M zinc sulphate
solution needed? [23 g] (Lit .: e.g. Röbisch p. 58)
12. 180 g of a solution with a mass fraction of B of 68% are diluted with 140 g of solvent.
What is the mass fraction of B in the diluted solution? [ω = 38.25%] (Lit .: e.g. Röbisch p. 63)
13. A hydrochloric acid with a mass fraction of HCl of 26% should be diluted with water to produce a
Hydrochloric acid of the mass fraction of HCl of 1% can be produced. What mass of water is for
To use thinning? [the mass of the starting solution must be increased by adding water
can be brought 26 times; e.g. 250 g water + 10 g HCl (ω = 26%)
14. From 0.9 kg of sulfuric acid with a mass fraction of 78%, the addition of conc. Sulfur-
acid (mass fraction: 98%) an 83% sulfuric acid can be produced. What mass of conc.
Sulfuric acid should be added? [m (H2SO4 conc.= 300 g] (Lit .: e.g. Röbisch p. 65)
15. Calculate the mass fractions of the elements Na, S and O in the compound Na2SO4!
[ωN / A= 32.38%, ωS.= 22.57%, ωO= 45.05%] (Lit .: e.g. Röbisch p. 67)
16. What is the maximum mass of pig iron that can be obtained from 2 tons of ore that contains a mass fraction of Fe3O4
of 72% has? [mFe= 1.04 t] (Lit .: e.g. Röbisch p. 68)
17. What mass of “crystal soda”, Na2CO3 • 10 H2O, is required to produce 750 g of 5%
Sodium carbonate solution must be weighed in? [101.2 g]
18. The elemental analysis showed the following mass fractions: carbon 75.88%, hydrogen 6.42%,
Nitrogen 17.81%. What is the ratio formula? [C5H5N] (Lit .: e.g. Röbisch p. 69)
19. How many g of lead (IV) oxide are produced in the oxidation of 20 g of lead? [23.08 g]
20. What mass of oxygen is consumed if 2 g of phosphorus are formed with the formation of P2O5 burned
become? [m = 2.58 g] (Lit .: e.g. Röbisch p. 72)
21. Hydrogen sulfide and sulfur dioxide should be converted into sulfur and water. Which
The mass of sulfur is obtained when 70 kg of H2S and 142 kg SO2 introduced into the reaction vessel
become? [m (S) = 98.7 kg] (Lit .: e.g. Röbisch p. 73)
22. From 2.38 g of a nickel (II) chloride hydrate, 1.3 g of the anhydrous yellow NiCl are obtained by heating2.
What was the composition (how many mol of water of crystallization) did the hydrate have?
[NiCl2 • .6 H2O; from M. Binnewies et al .: Exercise book General Chemistry, Spectrum, 2007]
23. A so-called physiological saline solution (corresponds to the salt content and thus the osmotic
Pressure of the blood plasma) has a molar concentration of c (Na+) = c (Cl-) = 154 mmol • l-1.
What is the mass of pure sodium chloride required to make 100 l of such a solution?
[900 g; from M. Binnewies et al .: Exercise book General Chemistry, Spectrum, 2007]
Lecture 3 (28.10.2010)
24. What is the law of the conservation of mass? Which two scientists discovered it?
25. The combustion of carbon to carbon dioxide is highly exothermic. The enthalpy of reaction is
394 kJ per mole of CO2. Calculate the corresponding loss of substance per mol using the EINSTEIN equation!
1 J = 1 Nm = 1 kg • m2 • s-2 [4,38 10-9 G]
26. What are the law of constant proportions and the law of multiple proportions?
27. Name the 4 atomic hypotheses of DALTON!
28. What are cathode rays and what are canal rays? How do they come about and what are their properties?
29. How was the neutron discovered (CHADWICK 1932; state the nuclear reaction equation!)?
30. a) How many iron atoms does a spherical pin head with a diameter of 1 mm contain?
ρ = 7.86 g • cm-3 r (Fe atom) = 126 pm
How long would the chain of iron atoms be if all atoms could be lined up while maintaining the Fe-Fe distance?
[4,4 1019; 1,11 107 km; from M. Binnewies et al .: Exercise book General Chemistry, Spectrum, 2007]
b) A square, thin gold foil (gold leaf) with an edge length of 10 cm has a mass of 0.01 g.
How thick is the gold foil? ρ (gold) = 19.5 g • cm-3
[5,18 10-5 mm; from M. Binnewies et al .: Exercise book General Chemistry, Spectrum, 2007]
Lecture 4 (29.10.2010)
31. Characterize the radiation that occurs during radioactive decay (3 types)!
32. What are the relationships between mass number, atomic number, neutron number, atomic number,
Number of protons, number of nucleons?
33. Write down the nuclear reaction equation for the α-decay of 226-Ra!
34. How do the mass number, number of neutrons and number of protons change in α-decay and in ß (-) decay?
35. Explain the terms "nuclides" and "isotopes" using two examples each!
36. Explain the age determination of C-containing organisms!
37. How is a chemical element characterized? What are pure elements and what are mixed elements?
Name each example!
38. Chlorine is made up of isotopes 35Cl and 37Cl. The rel. Atomic mass is 35.435. What is the ratio of both isotopes?
[78 % 35Cl, 22% 37Cl]
39. What is meant by the mass defect?
Lecture 5 (04.11.2010)
40. Formulate the nuclear reaction equation for the first artificial nuclear fission!
41. Describe RUTHERFORD's scattering experiment!
42. What are the basic ideas of RUTHERFORD's atomic model?
43. Formulate the nuclear reaction equation for the fusion of hydrogen and helium nuclei on the sun!
44. Formulate the reaction equations:
a) electrochemical decomposition of water
b) Representation of hydrogen from a base metal and dilute hydrochloric acid Winkler generator
c) Production of water gas / synthesis gas and conversion
d) Steam reforming process
45. Arrange the following substances according to their density, their diffusivity and their speed of propagation for sound:
Air, carbon dioxide, nitrogen, hydrogen, oxygen, helium!
Lecture 6 (05.11.2009)
46. How could one get dihydrogen from nasc. Differentiate hydrogen experimentally (state equation!)?
47. Sketch the energy curve during the chlorine-oxyhydrogen explosion! Formulate the corresponding reaction equations!
48. Explain why the dissociation of chlorine molecules (ΔH = + 243 kJ / mol) with green light (λ = 550 nm)
does not succeed (with calculation)! h = 6.62608 • 10-34 Js
49. Name the three groups of binary hydrogen compounds, the corresponding structural features and two examples each!
50. Write down an equation for each reaction
a) Ammonia synthesis (HABER-BOSCH process)
d) FISCHER-TROPSCH synthesis
e) Production of pure metals with dihydrogen!
Lecture 7 (11.11.2009)
51. Which contradiction made a further development of the atomic models according to DALTON and RUTHERFORD necessary?
How can the glow of gases in Geisler-Rören be explained?
52. How do emission spectra differ from absorption spectra?
53. What is meant by a spectrum?
54. What fundamental contradiction to the laws of classical electrodynamics gave rise to the formulation of the
two Bohr posulates?
55. What are Bohr's postulates? Also give an equation for each!
56. What is the basic idea of quantum theory (PLANCK 1900)?
57. Calculate the wavelength of a photon of frequency 1.2 x 1015 Hz! How big is the energy of such a photon? How large is
the energy of one mole of such photons (in kJ / mol)? What do you call this radiation? [λ = 250 nm; 7.9 x 10-19 J; 477 kJ; UV]
58. What is the energy associated with Bohr's fourth orbit, if the energy associated with the first orbit
-13.60 eV? What is the radius of this path (1st Bohr radius: 0.053 nm) [-0.85 V; 0.848 nm]
59. Why do excited atoms not provide a continuous spectrum, but a line spectrum?
Why is it characteristic of a certain element?
60. Describe the emission spectrum of hydrogen! How does it come about? With which equation can it be interpreted?
Explain the relationship between energy, frequency and wavelength of the observed radiation!
Lecture 8 (12.11.2009)
61. Name the disadvantages of Bohr's atomic model!
62. Which quantum numbers are there (name and symbol)? What is characterized by the respective quantum number? Which selection rules apply
63. Name the three rules / principles on which the structure of multi-electron systems and thus the PSE are based!
64. In which order are the energy levels (1s, 2s ...) occupied?
65. Give the electron configuration of 6C, 13Al, 23V, 24Cr, 30Zn,42Mon, 47Ag, 57La, 79Au and 92U on!
66. Give the valence electron configuration of 9F, 20Ca and 33As on!
67. Draw the energy level schemes (with sub-levels) for the elements contained in exercise 65!
Enter the quantum numbers for the electrons that were added last!
68. Characterize the 4 element blocks of the PSE with regard to their typical electron configuration!
Periodic tables on the Internet:
Lecture 9 (18.11.2010)
69. Define the following terms! Which trends are there within the main group from top to bottom and within
the period from left to right? Give a reason for each!
a) atomic radius
b) ionic radius
c) First ionization energy
d) electron affinity
70. Name the main components of the air with their respective proportions!
71. If you let a beaker with liquid air stand, the liquid turns blue after a while.
Give reasons for this phenomenon!
72. Name 3 groups of oxygen compounds with two examples each!
73. Name the three isotopes of oxygen!
74. Compare the amount of oxygen in the lithosphere and in the hydrosphere!
75. Compare the proportions of oxygen and nitrogen in the air and in the water!
76. What is the relationship between the solubility of a gas and the pressure and temperature?
Lecture 10 (19.11.2010)
77. Formulate the reaction equations for the oxidation of sulfur bloom and iron powder!
78. Explain the principle of the LINDE process!
79. What is catalysis and what is a catalyst? What is the effect of a catalyst based on?
Sketch the change in energy as a function of the course of the reaction for an endothermic reaction with and without
80. Write down the equations for the reaction
a) catalytic decomposition of hydrogen peroxide
b) thermal decomposition of potassium chlorate
c) Oxidation of acetylene
81. How many g of oxygen are needed to burn 1.5 g of hydrogen? How many g of water are created [12 g; 13.5 g]
82. As breathing gas for divers, O2/ He mixtures used. A mixture of 40.0 g O2 and 40.0 g of He has the required
Total pressure of 120 kPa. What are the partial pressures and the mole fractions of O?2 and hey?
Molenbruch: χA. = nA./ (nA. + nB.)
83. Can you name the main stages in the formation of oxygen on earth? How do you know that the earth's shell is in front of you
3.7 billion years ago did not contain significant amounts of oxygen?
84. Describe the structure of the ozone molecule! Also enter the resonance structures with the corresponding LEWIS
85. Formulate the reaction equations for the representation of ozone in the SIEMENS ozonizer!
How can one detect the ozone formed? (State the reaction equation!)
86. What is ozonolysis? What can it be used for?
87. What are three uses for ozone?
Lecture 11 (25.11.2010)
88. In the ozone layer, ozone is constantly being formed and then decays again. To do this, formulate the appropriate
89. Explain the destructive effect of chlorofluorocarbons on the ozone layer on the basis of reaction equations!
90. Why is ozone pollution often higher in some rural areas than in neighboring urban areas (equations
For interested parties:
Ozone in the atmosphere:
Chemistry and environmental pollution:
91. What is HEISENBERG's uncertainty relation (formulation and equation)?
92. What does DE BROGLIE's concept of matter waves mean (formulation and equation)?
93. Why is the experiment by DAVISSON and GERMER (1926) proof of the wave character of electrons?
94. How does Schrödinger describe the electron? What is the special meaning of | ψ |2? Provide a phrase for that
95. What is meant by an atomic orbital? How do atomic orbitals differ?
Lecture 12 (26.11.2010)
96. Which three quantum numbers result from the Schrödinger equation? What do they each characterize?
Which selection rules apply to them?
97. Sketch the spatial shape of the s, the three p and the five d orbitals in a spatial coordinate system!
In the case of the p and d orbitals, also give the sign of the wave function in the respective orbital lobes!
98. Give the electron configuration of 25Mn, 32Ge and 38Sr in the basic state (PAULING notation)!
99. When do chemical bonds develop (consideration of energy and electron configuration)?
100. Give an overview of the classification of chemical bonds!
101. Explain the nature of the covalent, the ionic and the metallic bond using two examples each!
102. How is the nuclear distance (bond length) of 74 pm to be explained in the hydrogen molecule, although the 1st Bohr radius is 53 pm?
103. Explain the basic ideas of the LEWIS theory! Enter the LEWIS formulas of dihydrogen, dichloro, dioxygen, dinitrogen,
Carbon dioxide and water!
104. What is meant by the octet rule?
Lecture 13 (02.12.2010)
105. Formulate the mesomeric boundary structures (resonance formulas) for ENT3, NO3-, CO32- and C6H6!
106. Explain the principle of octet expansion using the PCl examples5 and SF6! What kind of hybridization is there in each case?
107. What is understood by a coordinative (dative) bond? Explain this using the examples [Co (NH3)6]3+ and [Cu (NH3)4(H2O)2]2+!
108. Explain the LEWIS acid-base concept using two examples you have chosen!
Lecture 14 (03.12.2010)
109. Calculate the ionic bond fraction in HCl (μ (HCl)exp = 3,4410-30 Cm, q = 1.6 • 10-19 C, d = 127 pm)!
110.Draw CO's LEWIS formulas2, HCl and H2O! Mark the partial loads! Why is CO2 a non-polar one
Molecule, as opposed to HCl and H2O?
111. How do dipole molecules align and what are the consequences?
112. Explain the relationship between strength of the covalent bond, dissociation energy, bond length and degree of bond
the examples difluorine, dioxygen and dinitrogen!
113. For HCl, the sum of the covalent radii of the two binding partners is 135 pm. Experimentally, however, a bond length of 127 pm is used
found. Give reasons for this finding!
114. What is the relationship between the multiple bond content and the bond length?
115. Name the 4 rules of the VSEPR model and explain each of them using an example!
116. What geometries are for molecular compounds of general composition
FROM2, AB3, FROM3E, AB4, AB4E.2, AB5 and AB6 to be expected?
Arrange the connections BeCl2, CO2, BF3, CO32-, NO3-, CH4, PCl5,
SF6, SO32-, NH3, XeF4, PCl3 the respective structure type!
Lecture 15 (09.12.2010)
117. Explain the basic idea of the VB theory using the examples of dihydrogen and dinitrogen.!
118. Compare σ and π bonds with each other! How do they come about? Also go for the free rotation
and the strength of the bonds!
119. What is meant by the double bond rule? How can it be justified? What condition must be met in order to cancel it?
120. Using the VB theory, name two criteria for strong ties!
121. What is meant by hybridization? Explain the bonding relationships in methane, ethane, ethene, ethine! Give it
also take a look at the characteristic bond angles!
122. Name 6 types of hybridization, the associated number of hybrid orbitals formed with their respective orientation (geometry)
and one example each!
123. What type of hybridization and what geometry is there?
N2, PCl5, SF6, C3H8, CO2, SiH4, BF3, C2H4, NH3, HgCl2, [AuCl4]-, [Cr (NH3)6}3+, H2O, NO3-, [PtCl4]-, SbCl5, [SiF6]2-, C2H2.
Lecture 16 (10.12.2010)
124. What is the basic idea of MO theory? What is the LCAO method?
125. Draw the MO scheme for dihydrogen! Why does he exist2 Not? Calculate the corresponding bond orders!
126. Oxygen is paramagnetic. What is meant by this property? What is the cause? Formulate the mesomers
(LEWIS) limit formulas that reflect this situation on the one hand and the binding order on the other!
127. Draw the MO scheme of dioxygen (triplet oxygen)! Explain the term triplet oxygen!
128. In which structural features (indicate π * orbitals) and in which properties does the triplet oxygen differ from the singlet oxygen
Oxygen? So what can singlet oxygen be used for?
129. Give reasons for the occurrence of redder Luminescence when chlorine is introduced into an alkaline H2O2-Solution!
130. How does the color of the leaves come about in autumn?
Lecture 17 (16.12.2010)
131. Describe the structure of the hydrogen peroxide molecule (111 °!) And give a reason for it!
How is the strength of the O-O bond to be assessed?
132. Write down the redox equations for the action of H2O2 (sour mileu)
a) as an oxidizing agent (iodide ions are oxidized to iodine) and
b) as a reducing agent (permanganate ions become manganese2+Ions reduced, with simultaneous O2-Education)!
133.Draw the state diagram of the water with the following characteristic curves or points:
Melting, sublimation, vapor pressure curve; Melting, triple and critical point.
Derive five essential conclusions from this!
134. How can the bond angle in the water molecule of 104.5 ° be justified?
135. Name six properties of water!
136. Which requirements have to be fulfilled so that hydrogen bonds can be formed?
How is the strength of the hydrogen bonds compared to van der Waals bonds and ionic bonds?
137. Describe the structure of ice!
Lecture 18 (17.12.2010)
138. What is meant by the density anomaly of water! Name four resulting consequences!
139. How is the energy used to overcome the lattice energy when releasing
a) Molecular compounds (molecular lattice, e.g. sugar)
b) Ionic compounds (ion lattice, e.g. NaCl)
140. Define the terms cations, anions, cathode, anode!
141. How do electrolytes differ from non-electrolytes (with justification, give two examples!)?
142. Write down the equations for the dissociation of sodium chloride, calcium chloride and ammonium nitrate!
143. The conductivity of the following systems was measured: pure water, sodium chloride solution, hydrochloric acid, sucrose solution, acetic acid.
Arrange the named compounds according to their increasing conductivity. How are the differences to be justified?
144. What are cryoscopy and ebullioscopy? What can the methods be used for?
145. Name 3 colligative properties! What is meant by this term?
146. The increase in boiling point and lowering of the freezing point of solutions of the same concentration were investigated. Explain
the finding that an NaCl solution has twice the effect compared to the solution of a non-electrolyte, a Na2SO4-Solution
even shows a tripled effect!
147. What is meant by osmosis? Describe and justify the changes in the erythrocytes on contact with
a) water, b) NaCl (ω = 20%) or c) (ω = 0.9%; isotonic saline solution)!
148. Why do ripe cherries or tomatoes burst when it rains?
For interested parties:
Osmosis power plant
Lecture 19 (06.01.2011)
149. What is the essence of the ionic bond? Characterize the following compounds in terms of their type of binding:
CO2, KCl, HCl, ZnS, O2, HI!
150. Name three properties that are characteristic of ionic compounds (salts)!
151. What is meant by lattice energy? What is your knowledge important for? What is the relationship between the charges?
the cations and anions as well as the distance and the lattice energy?
152. The ionic radii are particularly important for the structure of the ionic compounds. What is the relationship between
a) Ionic radius and coordination number
b) Size of the ionic radius of cations and anions
c) Ionic radius and position in the main group and within the period (only indicate tendency!)
d) Ionic radius of positive ions (same ion, same coordination number) and ion charge.
153. Name general principles for the structure of ionic compounds (symmetry, coordination number, packing)!
154. Name two structures for KA and two for KA2-Structures a lattice type!
For interested parties:
(AB grid: cesium chloride grid, sodium chloride grid, zinc blende grid
FROM2-Grid: fluorite grid * rutile grid * cristobalite grid *) * not relevant to the test
155. State the coordination number, geometry and radius ratio for every two ion lattices you have chosen!
156. What is HESS's theorem?
157. Explain how the lattice energy can be determined with the help of the HABER-BORN cycle!
Lecture 20 (07.01.2011)
158. What are the two main laws of thermodynamics?
159. Many salts are readily soluble in water despite their high lattice energy? How can this fact be justified?
160. Discuss the influence of energy and probability parameters on the voluntary course of chemical reactions!
161. What is meant by the activity of an ion in a solution? In which case one can count on the concentration instead of the activity
be (with justification)?
162. Calculate K for the synthesis of hydrogen iodidepif Kc at 490 ° C is 45.9!
163. To what fraction does hydrogen iodide decompose?
at 25 ° C, when the equilibrium constant has the value 808! [6.57%; Lit: e.g. Uhlemann et al. CfL vol. 16, p. 20]
164. In water-gas equilibrium
the following parts by volume are present at 530 ° C: CO2 and H2 at 33.33% each, CO and H2O at 16.66% each.
Calculate the equilibrium constant. What are the concentration ratios at 970 ° C when
the equilibrium constant for this temperature is 0.64? [4; CO2 and H2 22.1% each, CO and H2O at 27.8% each; Lit. see above]
165. When 1 mol of ethanol reacts with 1 mol of acetic acid, 2/3 mol of ethyl acetate and 2/3 mol of water are formed.
Calculate the equilibrium constant! [4; Lit. see above]
166. Discuss the position of the chemical equilibrium as a function of the equilibrium constant K and the relationship
between free enthalpy of reaction and K!
167. Explain the influence of the concentration of the reaction participants on the position of the chemical equilibrium using the example of oxidation
of sulfur dioxide!
168. Explain the influence of temperature and pressure (total pressure) on the position of the chemical equilibrium using the example
of ammonia synthesis (ΔH = -92.4 kJ / mol)!
Lecture 21 (13.01.2011)
169. What is meant by the principle of Le Chatelier and Braun?
170. Discuss the influence of temperature and concentration (addition of water) on the position of the following equilibrium.
The forward reaction is endothermic.
171. Explain using the example of hydrogen iodide synthesis
a) the relationship between the equilibrium constant and the rate constants of the forward and backward reactions
b) how the reaction rate of the forward reaction and the reaction rate of the reverse reaction is defined!
172. Which empirical rule applies to the influence of a temperature increase on the reaction rate?
173. Write down the Arrhenius equation in logarithmic form and in the form of a straight line equation! Interpret the equation
with the help of the impact theory! Sketch the course of the function graphically! How can the activation energy be determined from this?
174. What is meant by catalysis? What is the effect of a catalyst due to?
Lecture 22 (14.01.2011)
175. Name five easily soluble and five poorly soluble substances (solvent: water)!
176. What is meant by the solubility product? Explain it using an example of your choice!
177. Calculate the solubility of
a) Mercury (II) sulfide, KL. = 310-54 mol2/ l2
b) lead (II) chloride, KL. = 2,1210-5 mol3/ l3 !
178. Explain why the titration end point in the argentometric chloride determination according to MOHR
can be recognized by the brown color of the solution (equations, solubilities)!
KL.(AgCl) = 2 • 10-10 mol2/ l2, KL. (Ag2CrO4) = 410-12 mol3/ l3
179. How do equi-ionic additives affect the solubility? By how much does the solubility of AgCl decrease if instead
of pure water a solution is used whose concentration of chloride ions 10-1 is (KL. so.)?
180. In which cases does the solubility improve with additives of the same ion! Formulate two reaction equations
as an an example!
181. How do foreign ionic additives affect the solubility? How is this behavior to be justified?
182. Explain why the solubility of barium sulfate improves 2.2 times when instead of water
a 0.01 M Mg (NO3)2-Solution is used! * Calculation * not relevant for the exam
Temperature dependence of the solubility
Lecture 23 (20.01.2010)
183. How did Lavoisier define acids? Name two disadvantages of this theory that made further development necessary!
184. What is meant by Liebig's "hydrogen theory" of acids?
185. How did Arrhenius define acids and bases? Name three disadvantages of the ARRHENIUS acid-base concept!
186. Explain the acid-base concept of BRÖNSTED and LOWRY!
187. Assign the following compounds or ions to the substance groups neutral acids, cationic acids, anionic acids,
Neutral bases, cation bases, anion bases, ampholytes to:
Water, phosphoric acid, hydrogen carbonate ions, carbonate ions, ammonia, hydroxide ions, acetate ions, acetic acid,
Hexaqua aluminum ions, hydrochloric acid, sulfate ions.
188. What is meant by the ionic product of water?
189. The pH of an orange juice was measured to be 2.4. A borax solution had a pH of 9.2.
Calculate the pOH values and the concentrations of hydrated hydrogen and hydroxide ions for both solutions!
Complete the following table!
Lecture 24 (21.01.2011)
190. Name two very strong, strong, medium and weak acids each!
191. What is meant by the leveling effect of water?
192. What reaction do you expect for the aqueous solutions of
d) [Fe (H.2O)6]3+
e) Well2SO4 ?
193. Calculate the pH values of
a) 0.5 M HCl
b) 0.05 M H2SO4
c) 0.2 M Hac; pKS.(Hac) = 4.74
d) 1 M NH4Cl; pKB.(NH3) = 4,75
e) sodium dihydrogen phosphate solution
(pKS1= 2.16; pKS2= 7.21; pKS3= 12.32 of phosphoric acid)
f) 0.2 M NaOH
g) 0.5 M Ca (OH)2 !
Lecture 25 (28.01.2011)
194. Decide which of the following reactions actually take place! Justify your decision!
a) hydrochloric acid + ammonia
b) ammonium chloride + sodium hydroxide solution
c) sodium hydrogen carbonate + ammonia
d) sodium chloride + acetic acid
pKS.(HCl) = -7; pKS.(NH4+) = 9.25; pKB.(OH-) = -1.74; pKS.(HCO3-) = 10,33
195. What is meant by the transition point of an indicator? Why is the transshipment area relevant for practical use?
Which structural feature must an indicator have?
196. Sketch the titration curves pH = f (vNaOH) for the titration of 0.1 M HCl or 0.1 M Hac with 1 M NaOH as standard solution!
Mark the equivalence points and the pKS.-Value of Hac! Select a suitable indicator for the respective titration
off (give reason)!
Areas of change for some indicators: thymol blue 1.2-2.8; Litmus 5.0-8.0; Phenolphthalein 8.2-10.0.
197. Explain the mode of operation of the buffer systems mentioned below using the respective reaction equations!
a) acetic acid / acetate buffer b) ammonia / ammonium chloride buffer c) dihydrogen phosphate / hydrogen phosphate buffer
d) carbonate / hydrogen carbonate buffer
How to proceed to prepare 1 l of buffer solution when starting solutions with a molar concentration of 2 mol / l are available
stand ? [355 ml Hac + 645 ml Naac; 849 ml NH4Cl + 151 ml NH3(aq)]
For interested parties:
Lecture 26 (28.01.2011)
199. How do complexes arise and how are they put together? Why do the subgroup elements form a particularly large number of complex compounds?
200. Formulate two complex formation reactions in each case, in which a) the solubility and b) the color change!
201. Define the terms coordination number and toothiness!
202. What is masking in complex chemistry? Explain this with an example!
203. Explain by giving the structural formulas and the adhesive atoms why ethylenediamine is a bidentate and ethylenediaminetetraacetate (EDTA)
is a hexadentate ligand! What are complexes with such multidentate ligands called and what are their distinguishing features?
204. Formulate the reaction equations! Also give the names of the resulting complexes!
a) Dissolution of a silver chloride precipitate by means of excess chloride ions
b) Dissolving a silver chloride precipitate with aqueous ammonia solution
c) Dissolving anhydrous nickel (II) chloride with complex formation
d) Reaction of the complex formed in c) with ammonia.
Lecture 27 (03.02.2011)
205. Give the names or formulas for the following complex compounds:
a) Hexaquachrome (III) chloride
b) Well3[(Cu (CN)4]
c) dichlorotetraaquachrome (III) chloride
d) Sodium Hexahydroxostannate (IV)
e) Potassium hexacyanoferrate (II) - yellow blood liquor salt
f) Potassium hexacyanoferrate (III) - red blood liquor salt
g) [Cu (NH3)4]2+
206. What is the meaning of complex compounds
a) in biological systems
b) in analytics
c) in technology?
207. What is meant by oxidations, reductions, oxidizing agents, reducing agents? Write down the equations for the reaction
a) Combustion of magnesium in the air
b) Reaction of magnesium with carbon dioxide
c) for the reaction of carbon monoxide with iron (III) oxide (blast furnace process)!
Assign the terms mentioned!
208. Which of the following reactions take place (complete the reaction equations!)?
Give reasons for the facts! Use the terms reducing power / oxidizing power, base / noble and high pressure to dissolve / tendency
a) Cu (s) + Ag+ (aq)
b) Zn (s) + Cu2+ (aq)
c) Cu (s) + Zn2+ (aq)
Lecture 28 (04.02.2011)
209. Assign the halide ions Cl-, Br- and I.- after increasing reducing power!
Which of the aqueous solutions of chlorine, bromine and iodine has the strongest oxidizing power?
210. State the oxidation numbers: phosphoric acid, sulfite ion, ozone, sodium peroxide, propanol, acetaldehyde, sodium hydride, sulfuric acid!
211. Write down the reaction equations
a) Iron nail in a copper (II) sulfate solution
b) hydrochloric acid + permanganate
c) oxalic acid + permanganate
d) Zinc + copper (II) ions (Daniell element)
e) hydrogen peroxide + permanganate
212. A sheet of copper is immersed in a silver nitrate solution. Explain the processes that lead to the formation of the electrochemical double layer
to lead !
213. Sketch an experimental arrangement in which the combination of the spatially separated corresponding redox pairs Cu (s) / Cu2+(aq)
and Ag (s) / Ag+(aq) can be used to generate electricity! Identify oxidation, reduction, anode and cathode,
Places with high and low electron pressure, places with higher and lower potential as well as the direction of the current flow!
214 How is a normal hydrogen electrode constructed and what is it used for?
215. Draw the standard electrode potentials (standard redox potentials, normal potentials) for the following in an energy diagram
corresponding redox couples!
E.0 (Al / Al3+) = -1.68 V, E.0 (Cu / Cu2+) = +0.34 V, E.0 (Ag / Ag+) = +0.80 V, E.0 (Zn / Zn2+) = -0.76 V.
What is the maximum potential difference that can be used in these examples to convert chemical energy into electrical energy
be (under standard conditions)?
Lecture 29 (10.02.2011)
216. Why do zinc and magnesium dissolve in hydrochloric acid while copper does not? Why is nitric acid suitable for dissolving copper?
Formulate the corresponding reaction equations!
217. Write down the reaction equations
a) Iron nail in a copper (II) sulfate solution
b) hydrochloric acid + permanganate
c) oxalic acid + permanganate
d) Zinc + copper (II) ions (Daniell element)
e) hydrogen peroxide + permanganate
f) lead battery!
218. How can you predict whether redox reactions will actually take place?
219. What voltage can be generated when combining two half-cells, in which the zinc electrode in zinc sulfate solutions of the
Amount of substance concentration 0.2 mol / l or 2 mol / l immersed? Temperature: 25 ° C; E.0 (Zn / Zn2+) = -0.76 V. Note: set c (Zn) = 1 mol / l!
220. What happens when chlorine is introduced into a solution that contains fluoride and bromide ions?
E.0 (2F-/ F2) = +2.87 V, E.0 (2Cl- / Cl2) = +1.36 V, E.0 (2Br-) / Br2) = +1.07 V. Explain your answer!
221. Decide whether the following reactions will take place (give reasons)! If necessary, formulate the corresponding reaction equations!
a) zinc nitrate + lead E.0 (Zn / Zn2+) = - 0.76 V, E.0 (Pb / Pb2+) = - 0.13 V
b) iron (II) chloride + copper E.0 (Fe / Fe2+) = - 0.41 V, E.0 (Cu / Cu2+) = + 0.34 V.
c) chloride + permanganate E.0 (Mn2+/MnO4-) = +1.51 V, E.0 (2Cl- / Cl2) = +1.36 V.
222. Why can chloride ions only from a strongly acidic KMnO4-Solution (e.g. pH = 0) can be oxidized to chlorine during the oxidation
from bromide to bromine in acetic acid solution (e.g. pH = 3) and the oxidation of iodide to iodine even succeeds from neutral solution.
E.0 (2Cl- / Cl2) = +1.36 V, E.0 (2Br-) / Br2) = +1.07 V, E.0 (2I- / I2) = +0.54 V, E.0 (Mn2+/MnO4-) = +1.51 V.
223. a) How can an Ag sheet be copper-plated?
b) Two Pb electrodes, to which a direct voltage is applied, are immersed in a lead (II) sulfate solution. What is happening?
c) "Zimmervulkan" (not relevant for the exam):
(NH4)2Cr2O7 -> 2 NH3 + CrO3 + H2O 2 NH3 + 2 CrO3 -> N2 + Cr2O3 + 3 H.2O total: (NH4)2Cr2O7 -> N2 + 4 H.2O + Cr2O3
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