copper

Examples of copper in the following topics:

• Copper

  • Hydrogen sulfides and sulfides react with copper to form various copper sulfides on the surface.
  • Hydrogen sulfides and sulfides react with copper to form various copper sulfides on the surface.
  • Copper(II) chloride and copper combine to form copper(I) chloride.
  • Among the numerous copper sulfides, important examples include copper(I) sulfide and copper(II) sulfide.
  • Many other oxyanions form complexes: these include copper(II) acetate, copper(II) nitrate, and copper(II) carbonate.

• Iron, Cobalt, Copper, Nickel, and Zinc

  • Copper is a member of a family of metals known as the "coinage metals," which includes copper, silver, gold, and roentgenium.
  • However, pure copper is too soft to have structural value, but copper alloys with zinc and tin to form harder brasses and bronzes.
  • Copper is easy to identify due to its reddish color.
  • Copper oxidizes—with some difficulty—to the +1 state in halides and an oxide, and to the +2 state in salts such as copper sulfate CuSO4.
  • Copper is heavily used due in many areas, including copper piping.

• Predicting the Products of Electrolysis

  • We take two copper electrodes and place them into a solution of blue copper sulfate (CuSO4) and then turn the current on.
  • We notice that the the initial blue color of the solution remains unchanged, but it appears that copper has been deposited on one of the electrodes but dissolved on the other.
  • At the positive anode, copper metal is oxidized to form Cu2+ ions.
  • This is why it appears that the copper has dissolved from the electrode.
  • Two copper electrodes are placed in a solution of blue copper sulfate and are connected to a source of electrical current.

• Predicting if a Metal Will Dissolve in Acid

  • The classic example is of zinc displacing copper:
  • Here, zinc is more active than copper because it can replace copper in solution.
  • If you immerse a piece of metallic zinc in a solution of copper sulfate, the surface of the zinc quickly becomes covered with a coating of finely divided copper.
  • The blue color of the solution diminishes as copper(II) ion is being replaced.
  • Consider, for example, the oxidation of copper by metallic zinc mentioned above.

• Voltaic Cells

  • In a typical voltaic cell, the redox pair is copper and zinc, represented in the following half-cell reactions:
  • Copper readily oxidizes zinc; the anode is zinc and the cathode is copper.
  • The zinc electrode produces two electrons as it is oxidized ($Zn \rightarrow Zn^{2+} + 2e^-$), which travel through the wire to the copper cathode.
  • The electrons then find the Cu2+ in solution and the copper is reduced to copper metal ($Cu^{2+} + 2e^- \rightarrow Cu$).
  • During the reaction, the zinc electrode will be used and the metal will shrink in size, while the copper electrode will become larger due to the deposited Cu that is being produced.

• Occurrence of Metals

  • The gold group consists of gold, copper, lead, aluminium, mercury, and silver.
  • Only gold, silver, copper and the platinum metals occur in nature in larger amounts.
  • So while copper and iron were known well before the Copper Age and Iron Age, they would not have a large impact on humankind until the technology to smelt them from their ores, and thus mass-produce them, appeared.
  • Other significant metallic alloys are those of aluminium, titanium, copper, and magnesium.
  • Chuquicamata, Chile, is the site of the largest circumference and second deepest open pit copper mine in the world.

• Equilibrium Constant and Cell Potential

  • This allows SO42- ions to flow freely between the copper and zinc solutions.
  • In order to calculate the standard potential, we have to look up the half-reactions of copper and zinc.
  • The polarity of the cell is determined by knowing that zinc metal is a stronger reducing agent than copper metal.
  • Therefore, the standard reduction potential for zinc is more negative than that of copper.
  • Thus, zinc metal will lose electrons to copper ions and develop a positive electrical charge.

• Naming Ionic Compounds

  • Its systematic name is copper (II) chloride, where copper's oxidation number is indicated in parentheses.
  • For example, Cu(NO3)2 is copper (II) nitrate, because the charge of two nitrate ions (NO3−1) is 2\cdot(-1) = -2.
  • This compound is therefore, copper (II) nitrate.
  • In the cases of iron and copper, the Latin names of the elements are used (ferrous/ferric, cuprous/cupric).

• Predicting Precipitation Reactions

  • Imagine you have two test tubes and three solutions: copper (II) chloride (CuCl2) solution, sodium carbonate (Na2CO3) solution, and sodium sulfate (Na2SO4) solution.
  • You put 5 mL of copper (II) chloride into tubes 1 and 2.
  • You can automatically exclude the reactions where sodium carbonate and copper (II) chloride are the products because these were the initial reactants.
  • You know that sodium chloride (NaCl) is soluble in water, so the remaining product (copper carbonate) must be the one that is insoluble.
  • Once again, the reactions where sodium sulfate and copper (II) chloride are the products can be excluded, since they were the initial reactants.

• Chelating Agents

  • Consider the two equilibria, in aqueous solution, between the copper(II) ion (Cu2+) and ethylenediamine (en) on the one hand and methylamine (CH3NH2 (MeNH2)) on the other.
  • In (1), the bidentate ligand ethylenediamine forms a chelate complex with the copper ion.
  • Under conditions of equal copper concentrations and when the concentration of methylamine is twice the concentration of ethylenediamine, the concentration of the complex in (1) will be greater than the concentration of the complex in (2).