# cu + hcl reaction

Paul Flowers (University of North Carolina - Pembroke), Klaus Theopold (University of Delaware) and Richard Langley (Stephen F. Austin State University) with contributing authors. The electrode chosen as the zero is shown in Figure 17.4.1 and is called the standard hydrogen electrode (SHE). Its main significance is that it established the zero for standard reduction potentials. The Reaction of Magnesium with Hydrochloric Acid In this experiment you will determine the volume of the hydrogen gas that is produced when a sample of magnesium reacts with hydrochloric acid. Enter either the number of moles or weight for one of the compounds to compute the rest. Reaction 3 is observed because nickel is higher up on the activity series of metal than copper. Limiting reagent can be computed for a balanced equation by entering the number of moles or weight for all reagents. We also acknowledge previous National Science Foundation support under grant numbers 1246120, 1525057, and 1413739. Copper does not react with HCl because HCl is not an oxidising acid. Reversing the reaction at the anode (to show the oxidation) but not its standard reduction potential gives: \begin{align*} In cell notation, the reaction is, \[\ce{Pt}(s)│\ce{H2}(g,\:1\: \ce{atm})│\ce{H+}(aq,\:1\:M)║\ce{Cu^2+}(aq,\:1\:M)│\ce{Cu}(s), Electrons flow from the anode to the cathode. Chemical reactions tend to involve the motion of electrons, leading to the formation and breaking of chemical bonds.There are several different types of chemical reactions and more than one way of classifying them. A galvanic cell consisting of a SHE and Cu2+/Cu half-cell can be used to determine the standard reduction potential for Cu2+ (Figure $$\PageIndex{2}$$). Reaction stoichiometry could be computed for a balanced equation. A chemical reaction is a process generally characterized by a chemical change in which the starting materials (reactants) are different from the products. *Response times vary by subject and question complexity. In many cases a complete equation will be suggested. I might come back with some new questions, but for now, thanks. The standard reduction potential can be determined by subtracting the standard reduction potential for the reaction occurring at the anode from the standard reduction potential for the reaction occurring at the cathode. Given the following list of substances and the common reaction templates answer the questions below: NaOH H2 C8H18 CaCO3 Zn H2SO4 O2 Cu(NO3)2 acid + base ----> water + ionic compound metal + oxygen -- … $\ce{Mg}(s)+\ce{2Ag+}(aq)⟶\ce{Mg^2+}(aq)+\ce{2Ag}(s) \hspace{20px} E^\circ_\ce{cell}=\mathrm{0.7996\: V−(−2.372\: V)=3.172\: V} Reaction of copper immersed in HCl. (s)Cu. The voltage is defined as zero for all temperatures. The volume of the hydrogen gas produced will be measured at room temperature and pressure. Galvanic cells have positive cell potentials, and all the reduction reactions are reversible. What is the balanced equation of copper metal and silver nitrate? 1)How can I tell if a reaction like Zn + Hcl -> ZnCl2 + H2 can happen or not? Copper is a very unreactive metal, and it does not react with hydrochloric acid. Textbook content produced by OpenStax College is licensed under a Creative Commons Attribution License 4.0 license. It is single replacement. When calculating the standard cell potential, the standard reduction potentials are not scaled by the stoichiometric coefficients in the balanced overall equation. The reduction reactions are reversible, so standard cell potentials can be calculated by subtracting the standard reduction potential for the reaction at the anode from the standard reduction for the reaction at the cathode. As the name implies, standard reduction potentials use standard states (1 bar or 1 atm for gases; 1 M for solutes, often at 298.15 K) and are written as reductions (where electrons appear on the left side of the equation). Platinum, which is inert to the action of the 1 M HCl, is used as the electrode. Consider the cell shown in Figure $$\PageIndex{2}$$, where, \[\ce{Pt}(s)│\ce{H2}(g,\:1\: \ce{atm})│\ce{H+}(aq,\: 1\:M)║\ce{Ag+}(aq,\: 1\:M)│\ce{Ag}(s)$, Electrons flow from left to right, and the reactions are. [ "article:topic", "Author tag:OpenStax", "standard cell potential", "standard hydrogen electrode", "standard reduction potential", "authorname:openstax", "showtoc:no", "license:ccby", "transcluded:yes", "source[1]-chem-38305" ], https://chem.libretexts.org/@app/auth/2/login?returnto=https%3A%2F%2Fchem.libretexts.org%2FCourses%2FBrevard_College%2FCHE_104%253A_Principles_of_Chemistry_II%2F01%253A_Electrochemistry%2F1.07%253A_Standard_Reduction_Potentials, $\mathrm{+0.80\: V}=E^\circ_{\ce{Ag+/Ag}}−E^\circ_{\ce{H+/H2}}=E^\circ_{\ce{Ag+/Ag}}−0=E^\circ_{\ce{Ag+/Ag}}$, $E^\circ_\ce{cell}=E^\circ_\ce{cathode}−E^\circ_\ce{anode}=E^\circ_{\ce{Ag+/Ag}}−E^\circ_{\ce{Cu^2+/Cu}}=\mathrm{0.80\: V−0.34\: V=0.46\: V}$, $$\ce{3Ni}(s)+\ce{2Au^3+}(aq)⟶\ce{3Ni^2+}(aq)+\ce{2Au}(s)$$, $E^\circ_\ce{cell}=E^\circ_\ce{cathode}−E^\circ_\ce{anode}=\mathrm{1.498\: V−(−0.257\: V)=1.755\: V}$, 1.6: Batteries- Using Chemistry to Generate Electricity, 1.8: Electrolysis- Using Electricity to Do Chemistry. 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