Lab #5 lab sample


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Lab #5 lab sample

  1. 1. LAB # 5: Metals and Single Replacement ReactionsPurpose: Compare different metals and their reactions with hydrochloric acid. Construct a model of oxidation and reduction in single replacement reactions.Pre-lab Questions: 1. Describe the bond that holds together metals. Explain how this allows it to conduct electricity. Metallic bonds hold the metal atoms in fixed positions. These free moving electrons allow the metal to conduct electricity. The electrons have a charge and can move. 2. Metals are composed of atoms - explain why the metal and its atoms are not soluble in water. Metals are composed of atoms with no charge. These atoms with no charge are not attracted to polar water molecules. 3. Explain why metal ions are soluble- especially in ionic compounds with nitrate. What happens to the ions in solution? Draw a visual representation. Metal ions have a positive charge, so they are attracted to water molecules. When ions dissolve, the ions dissociate – split up – and form ion-dipole attractions to water molecules. The strength of the ionic bond determines the solubility of the metal in an ionic bond. When nitrate ions are the negative ions, the compound is always soluble. 4. Is sodium more stable as an atom or an ion? Compare the reaction of sodium atoms (in sodium metal) to sodium ions (in sodium chloride salt) when they are added to water. Sodium atoms were added to water in the form of solid sodium. The sodium bursts into flames and frequently explodes. The sodium ions in sodium chloride (table salt) added to water is quite boring. No explosions,; it just dissolves. The sodium ion must be more stable. They do not change much and release energy in the process of becoming more stable.Results: Data: Adding different metals to Hydrochloric Acid Metal Observations Ti Tf ∆T Zinc Magnesium Copper Aluminum Metal Observations Ti Tf ∆T Zinc Bubbling- zinc floats and sinks increase Magnesium Immediate bubbling and piece of Mg gone increase Copper No reaction No ∆ Aluminum No reaction No ∆Initial Analysis: 5. Compare the reactions of each metal. Rank the metals from most reactive to least. Mg Zn Al Cu reacted immediately react in 1 M reacted in the 6 M never reacted 6. Which reactions were noticeably exothermic? Support with observations. The zinc, magnesium and aluminum reacted and all felt hot (temperature increased). The increase in temperature shows an increase in kinetic energy. This kinetic energy must have come from the reaction. So the reactions are exothermic. 7. In all the reactions, what is one of the visible products? Support with observations. Bubbles were observed in all the reactions. This shows that a gas was produced. The molecules are far apart and neither attracted to the water nor each other.
  2. 2. 8. What happened to the atoms of the metals when the metals reacted? Where are they now? Are they still atoms? Explain your thinking. The metals were composed of atoms – had no charge and were not attracted to polar water. After the reaction, the metals must be in solution since we no longer see them. They must now have a charge to be attracted to the water. So the metals must now be ions. 9. Propose a hypothesis about the products of these reactions. What do think happened to the atoms? Where are they now? What do you think the gas is? We started with only metals and HCl. The metals that are now in the solution must have formed positive ions. There must also be negative ions- chloride ions since hydrogen ions are positive. The hydrogen must have become a gas. 10. Propose a method to capture and identify the products of the reaction.Demonstration: Record observations of aluminum in 6.0 M hydrochloric acid.Analysis of the Reactions: 11. EQUATIONS TO SUMMARIZE THE REACTIONS: Write a balanced equation to summarize the reaction of the metals that reacted with hydrochloric acid. Be sure to include the reactants and products and the symbols. Mg (s) + 2 HCl (aq)  MgCl2(aq) + H2(g) Zn (s) + 2 HCl (aq)  ZnCl2(aq) + H2(g) 2 Al (s) + 6 HCl (aq)  2 AlCl3(aq) + 3 H2(g) 12. REACTANTS: a. Describe the metals- What kind of particles? What charge on the particles? Compare the number of protons and electrons in atoms. Identify and describe the bond. Metals are composed of atoms with no charge, so the number of positive protons must equal the number of negative electrons. The atoms are held in fixed positions by metallic bonds – a sea of free moving electrons shared by all the atoms. b. Describe and draw the ions of hydrogen chloride in aqueous solution (hydrochloric acid- an ionic acid with hydrogen ions in solution). Include the numbers of protons and electrons in each ion. Acids (inorganic acids) have hydrogen ions and a negative ion. In the case of hydrochloric acid, the negative ion is chloride. Since hydrochloric acid is a strong acid – the ions completely dissociate in solution – that is all the HCl splits up into H+ ions and Cl- ions in solution. These ions are surrounded by water molecules. The hydrogen ions (H+) each have one proton (+1) and no electrons for a net charge of positive one. The chloride ions (Cl-) each have seventeen proton (+17) and eighteen electrons (–18) for a net charge of negative one. 13. PRODUCTS: a. Draw and describe the gas product produced in all of the reactions. o What kind of bond holds together the atoms? The two hydrogen atoms have the same pull on the electrons. They share the electrons equally – a nonpolar covalent bond. Since the electrons spend the same time around each atom, the ends have no oppositely charge. o How did the hydrogen ion change? What happened to the electrons? The hydrogen ions had no electrons. The hydrogen ions have become hydrogen atoms and they are now sharing a pair of electrons. They did not have any electrons, so they must have gained electrons. They were reduced. o Was the hydrogen oxidized or reduced? Write a half-reaction to show the change. 2H+ (aq) + 2e-  H2 (g)
  3. 3. b. The magnesium atom became an ion of magnesium in the reaction. o Draw and describe the magnesium ion in solution. What is the charge on the ion? Magnesium ions in solution have a charge of positive two (+2) and attracted to the water molecules by ion-dipole attractions. o How many protons and electrons in the atoms before the reaction? o How many protons and electrons in the ions after the reaction? Before After Mg (s)  Mg2+ (aq) Atoms Ions 12 p 12p 12 e- 10 e- 0 +2 o Did atom gain or lose an electron to become an ion? Magnesium had 12 electrons and ends with only 10 electrons, so it lost two electrons. Write half-reaction to show this change. Was the atom oxidized or reduced? Since magnesium lost electrons, it was oxidized Mg (s)  Mg2+ (aq) + 2e- need these two electrons to balance both sides c. The aluminum atom became an ion of aluminum in the reaction. Answer all the bullets in b. above for the aluminum (except that a drawing and description of the ions in solution is not necessary since it would be the same as for magnesium). Before After Al (s)  Al3+ (aq) Atoms Ions 12 p 13p 12 e- 10 e- 0 +3 OXIDATION HALF-REACTION Al (s)  Al3+ (aq) + 3e- d. The zinc atom became a +2 ion of zinc in the reaction. Zinc is a transition metal. Again. Answer all the bullets in b for zinc. Before After Zn (s)  Zn2+ (aq) Atoms Ions 30 p 30 p 30 e- 28 e- 0 +2 OXIDATION HALF-REACTION Zn (s)  Zn2+ (aq) + 2e- e. The copper does not react with hydrochloric acid, what does that suggest about copper? Copper must be more stable and less reactive than the other metals.14. ENTHALPY- ENERGY IN REACTIONS: In the reaction of magnesium with hydrochloric acid, the energy change should have been noticeable. a. Was energy transferred into or out of the system? Describe your observational evidence. Energy transferred out of the system. Since temperature increased – the surrounding gained kinetic energy, which must have come from the reaction.
  4. 4. b. Was the reaction endothermic or exothermic? What happened in terms of stability? Since energy was transferred out, the potential energy decreased, so it became more stable. c. Did the energy content (PE) increase or decrease? Which has more PE the products or reactants? Since energy was transferred out, the potential energy decreased. The reactants have more PE. d. Describe the reactions in terms of the energy added to break bonds and form bonds. Energy is required to break the bonds between the metal atoms in the solid as well as between the hydrogen ions and water molecules. In addition, energy is required to pull off the electrons, since the electrons are attracted to the positive nucleus. Energy is released as new bonds form. The magnesium ions form ion-dipole attractions to the water molecules. The electrons pulled off of magnesium form attractions to the nucleus of hydrogen. The hydrogen atoms form nonpolar covalent bonds between the atoms to form hydrogen molecules. Since these reactions were exothermic more energy must be released in the formation of the new attractions than was required to break the old attractions. The net change in energy is negative – or net transfer of energy is out. Draw an energy graph to summarize the reaction. Ein to Eout as break attractions form attractions H Net Δ H reactants negative products e. Is this reactions favored or not favored based upon the change in enthalpy? Explain Since the products have less potential energy they are more stable – so this reaction is favored to occur based upon enthalpy.15. ENTROPY: a. What happened to entropy in these reactions? Explain your reasoning. Zn (s) + 2 HCl (aq)  ZnCl2(aq) + H2(g) In the beginning, we have solid (fixed positions) and solution (free to move). In the end, we have a gas (free to move and far apart) and solution (free to move). The disorder increased since the particles have a greater freedom of movement in the end. With more possible location there is more disorder. b. Are these reactions favored or not favored based upon the change in entropy? Explain. Since the disorder increased, the reaction is favored based upon entropy. So, this is one of those reactions that is favored based upon both enthalpy and entropy.16. ATOMS COMPARED TO IONS. Differentiate between atoms and ions. Compare magnesium as an ion with magnesium as an atom. What is different? What is the same? Are the atoms soluble? Are the ions soluble? Which is more stable?17. REACTION RATES. a. Use a collision model to explain why the aluminum reacted faster with the more concentrated hydrochloric acid. See end of document for graphic showing different concentrations. In these reactions hydrogen ions in solution collide with the metals atoms. NET IONIC EQUATION H+ (aq) + Al (s)  Al3+ (aq) + H2 (g) REDUCTION HALF-REACTION 2H+ (aq) + 2e-  H2 (g)
  5. 5. OXIDATION HALF-REACTION Al (s)  Al3+ (aq) + 3e-The hydrogen ions must collide with the aluminum atoms with enough kinetic energy to break the attractions between atoms as well as the attraction of the negative electrons to the positive nucleus.If they collide with enough KE, an effective collision occurs and the reaction proceeds as new bonds form and release energy.In the more concentrated solution, there are more ions in the same amount of solution. More collisions will occur between the hydrogen ions in the solution and the metal.The reaction will be faster.b. Use a collision model to explain what would happen to the rate of reaction of zinc with hydrochloric acid if the temperature of the acid were decreased.Decreasing the temperature decreases the average kinetic energy per particle.With less KE the particles will collide less often and with less KE.Thus less collisions will be effective per second and the reaction will be slower.
  6. 6. Lab #5 Worksheet Name _______________________________ period ____Predicting the Products of Single Replacement ReactionsUse activity series (Holt textbook page 327) to determine whether the reaction occurs. If the reaction occurs,predict the products and balance the equation. Use the most common ion when more than one is possible. If thereaction does not occur, the write NR (no reaction) and explain why.1. magnesium and calcium chloride Mg (s) + CaCl2 (aq)  NR – Mg less active than Ca so does not replace2. lead with copper (II) sulfate Pb (s) + 2 CuSO4 (aq)  Pb(SO4)2 (s) + 2 Cu (s)3. aluminum with barium nitrate Al (s) + Ba(NO3)2 (aq)  N.R.4. zinc carbonate with sodium 2 Na (s) + ZnCO3 (aq)  Na2CO3 (aq) + Zn (s)5. lithium with sodium chloride Li (s) + NaCl (aq)  LiCl(aq) + Na (s)6. magnesium iodide with potassium MgI2 (aq) + 2 K (s)  2 KI(aq) + Mg (s)7. hydrochloric acid with silver HCl (aq) + Ag (s)  NR8. tin (IV) chloride with strontium 2 Sr (s) + SnCl4 (aq)  Sn(aq) + 2 SrCl2 (aq)9. zinc with copper (II) nitrate Zn (s) + Cu(NO3)2 (aq)  Zn(NO3)2 (aq) + Cu(s)10. lead with magnesium sulfate Pb (s) + MgSO4 (aq)  NR11. Sacrifical anodes are often used to protect a metal. The Alaskan pipeline is protected by a zinc cable thatrusts alongside it. Which metal will be oxidized first (which is more active and will therefore react first)?12. Galvanized iron nails are coated with zinc. How does this protect the nail from rusting?13. Select and defend a metal that might be used to protect tin cans from oxidizing.
  7. 7. 1.0 M H+ 3.0 M H+ Same volume More ions So – more concentrated