Percent Copper in Brass
Materials Included In Kit
Brass sample, 75 g
Copper(II) nitrate solution, Cu(NO3)2, 0.1 M, 75 mL
Copper(II) nitrate stock solution, Cu(NO3)2, 0.40 M, 200 mL
Copper(II) sulfate solution, CuSO4, 0.1 M, 75 mL
Iron(III) chloride solution, FeCl3, 0.1 M, 75 mL
Iron(III) nitrate solution, Fe(NO3)3, 0.1 M, 75 mL
Nitric acid, concentrated, HNO3, 15.8 M, 75 mL
Zinc nitrate solution, Zn(NO3)2, 0.1 M, 75 mL
Zinc sulfate solution, ZnSO4, 0.1 M, 75 mL
Pipets, serological, 10-mL, 12
Additional Materials Required
(for each lab group)
Water, distilled or deionized
Balance, 0.001-g precision, 1–2 to share
Beaker, 50-mL, with watch glass
Cuvets or test tubes, 13 x 100 mm, 7
Graduated cylinder, 50-mL
Hot plate (shared)
Kimwipes or lens tissue
Pipet bulb or pipet filler
Spectrophotometers, 3–6 to share
Test tube rack
Volumetric flask, 100-mL
Concentrated nitric acid is severely corrosive, a strong oxidizer and toxic by ingestion and inhalation. Keep sodium carbonate acid neutralizer on hand to clean up acid spills. Reactions of nitric acid with metals generate nitrogen dioxide, a toxic, reddish-brown gas. Work with nitric acid only in a fume hood. Copper(II) sulfate, copper(II) nitrate and zinc nitrate solutions are toxic and irritating to skin and body tissue. Iron(III) chloride and iron(III) nitrate solutions may be skin and body tissue irritants. Zinc sulfate is a skin irritant. Wear chemical splash goggles, chemical-resistant gloves and a chemical-resistant apron or lab coat. Please follow all normal laboratory safety guidelines. Remind students to wash their hands thoroughly with soap and water before leaving the laboratory. Please review current Safety Data Sheets for additional safety, handling and disposal information.
Please consult your current Flinn Scientific Catalog/Reference Manual for general guidelines and specific procedures, and review all federal, state and local regulations that may apply, before proceeding. Excess concentrated acid may be neutralized according to Flinn Suggested Disposal Method #24b. Copper(II) sulfate, copper(II) nitrate, iron(III) chloride, iron(III) nitrate, zinc nitrate and zinc sulfate solutions may be rinsed down the drain with excess water according to Flinn Suggested Disposal Method #26b. All spectroscopic sample solutions may be combined and neutralized according to Flinn Suggested Disposal Method #24b.
- Ideally, each lab group should have a spectrophotometer or colorimeter to measure the absorbance of their solutions. However, by first preparing all solutions, then measuring the absorbance values of one solution after another, it is possible to complete this lab with only a few of these instruments in the classroom.
- Before attempting this laboratory, students should be able to:
- Use basic laboratory instruments, including a balance, volumetric pipets and a spectrophotometer and colorimeter.
- Display data in graphic form.
- Balance an oxidation–reduction reaction.
- Assign the lab groups one of the following sets of salt solutions for the Introductory Activity.
0.1 M copper(II) nitrate
0.1 M iron(III) nitrate
0.1 M copper(II) sulfate
0.1 M iron(III) sulfate
- Volumetric flasks ensure the greatest accuracy in this quantitative analysis lab. If you do not have enough flasks for each student group, students may share by transferring their prepared solutions to a labeled bottle and then passing on their flasks (rinsed clean!) to the next group.
- Students are asked to determine a suitable wavelength at which to measure the absorbance of the brass solution. Discuss the factors that influence this choice—no interferences, adequate range of concentrations and sensitivity, highest absorbance value ≤1.00, etc. Students will use the 0.1 M Cu2+(aq) ion spectrum to determine the wavelength, but the calibration curve should extend to 0.4 M. They should select a top absorbance value near 1.000 (say 0.800) and divide this value by 4 (since 0.4 M is four times 0.1 M). The students should look at the spectrum to find the wavelength where the 0.1 M Cu2+(aq) ion has an absorbance value of 0.200.
- The reaction of brass with concentrated nitric acid completely dissolves the zinc and copper metals in the brass.
- In Part B, students are to add 15.8 M nitric acid to their samples to dissolve the brass. Ask students to calculate the amount needed to dissolve the brass, assuming the sample is all copper.
- The volume of HNO3(aq) that is required is calculated from the overall balanced equation:
Cu + 4HNO3 → Cu(NO3)2 + 2NO2 + 2H2O
The setup for the calculated minimum volume is: mL HNO3 needed =
1.0 g brass(as Cu) (1 mol Cu /63.55 g Cu) (4 mol HNO3/1 mol Cu) (1000 mL HNO3/15.8 mol HNO3)
The minimum volume is 4 mL.
- The nitric acid–brass reaction is quite vigorous and produces both NO and NO2 gases. The NO also reacts with the oxygen in the air to produce the toxic reddish-brown gas NO2. A fume hood is required for this reaction. While the reaction is very rapid at first, as the acid is consumed, the rate slows. If the beaker is placed on a hot plate at a low setting, the metal should quickly dissolve within 20 minutes. Do not allow the students to remove their beakers from the fume hood until they have added the 30 mL of distilled water, washing down any drops on the bottom of the watch glass into the beaker.
- The brass supplied in this kit has a nominal copper content of 70%, with a range, depending on the particular production lot, of 68.5% to 71.5%.
- Visible spectroscopy is an important tool in determining the composition of metals, either in industry, such as alloying amounts, or as a forensics tool when analyzing bullets taken from a crime scene.
Answers to Prelab Questions
- Dissolving brass requires an oxidizing acid, such as concentrated nitric acid. Nitrogen dioxide is produced as a byproduct in this reaction. Write a balanced chemical equation for the reaction of copper metal with concentrated nitric acid to produce copper(II) nitrate, nitrogen dioxide and water.
Cu(s) + 4HNO3(aq) → Cu(NO3)2(aq) + 2NO2(g) + 2H2O(l)
- Nitrogen dioxide is a toxic, reddish-brown gas. What safety precautions are needed in this inquiry lab to protect against this hazard, as well as the hazards attributable to the use of concentrated acid?
Work with concentrated nitric acid in a fume hood only. Wear chemical splash goggles, heavy-duty nitrile gloves, and a chemical-resistant apron or lab coat, especially when measuring or transferring nitric acid. When the brown fumes have dissipated, carefully add about 30 mL of distilled or deionized water to dilute the acid prior to removing the beaker from the hood.
- Copper(II) ions appear blue in aqueous solution. This is the transmitted color. The wavelengths of light that are NOT absorbed give rise to the perceived or transmitted color of a substance. Based on the principle of complementary colors, which colors or wavelengths of light would you expect to be most strongly absorbed by Cu2+ ions?
Solutions of copper(II) ions have a strong absorbance at visible wavelengths above 600 nm, corresponding to absorption of yellow, orange and red light. Yellow is the complementary color of blue light. The maximum absorbance is actually in the near infrared region, that is, at wavelengths >700 nm.
- Spectroscopy measurements may be made in either percent transmittance (%T) or absorbance (A). Based on the mathematical relationship between absorbance and transmittance, A = −log T, explain why calibration curves of absorbance versus concentration may deviate from a straight line when A < 0.1 and when A > 1.
Absorbance measurements are most accurate in the range 0.1–1.0. Recall that a spectrophotometer detects the difference or ratio between the power of the incident and transmitted light. At very low absorbance values, the percent transmittance is high (for A = 0.1, % T = 80%). Since almost all of the incident light goes through, there is not a great difference between the incident power and the transmitted light for the instrument to detect. At very high absorbance values the percent transmittance is very low (for A = 1.0, % T = 10%). The instrument does not accurately detect low power of transmitted light.
Calibration Curve for Cu2+ Solutions
Analysis for Percent Copper in Brass
Determine the mass of Cu dissolved in the brass solution and use these values to calculate the mass percent of Cu in brass.
Mass of brass sample = 0.973 g
Volume of dissolved brass solution = 100 mL
Sample absorbance = 0.251 at 630 nm
Slope of Cu2+ calibration curve = 2.3459 M–1
- Copper concentration in brass solution = (Sample absorbance)/(Slope of Cu2+ calibration curve)
= 0.251/2.3459 M–1
= 0.107 M
- Moles copper in brass sample = (Volume of dissolved brass sample)(Copper concentration in brass solution)
= (0.107 M)(0.1 L) = 0.0107 moles
- Percent copper in brass sample = [(moles Cu)(Molar mass Cu)/grams sample)] x 100
Answers to Questions
Guided-Inquiry Design and Procedure
- For each salt solution, determine the species (cation and/or anion) that is responsible for the absorbance spectrum. Explain your reasoning.
Only copper(II) and iron(III) salt solutions absorbed visible light. The spectra of copper(II) nitrate and copper(II) sulfate were almost identical, suggesting that the copper(II) cation is responsible for the visible light absorption at wavelengths > 600 nm. Similarly, the spectra of iron(III) chloride and iron(III) nitrate showed very close overlap as well, consistent with the iron(III) cation absorbing visible light from 400−550 nm. Chloride, sulfate and nitrate anions do not contribute to the visible spectra of the transition metal salts.
- Do Zn2+ ions absorb visible light? Discuss the answer in terms of (a) the color and appearance of Zn2+ aqueous solutions and (b) the electronic structure of Zn2+ ions. Hint: See the Background section for information on the electronic transitions of transition metal ions.
The zinc salt solutions were colorless and did not show any absorption of visible light in the 400−700 nm region. Zinc ions have a valence electron configuration of 3d10 after loss of their 4s2 electrons. Absorption of visible light by colored transition metal ions is associated with d–d transitions. Since Zn2+ ions have filled d-orbitals, electrons cannot be promoted or excited to a higher energy, unfilled d-orbital.
- Identify a suitable wavelength for analysis of Cu2+ ions in aqueous solution. The radiant intensity of light is highest in the middle of the visible range, and falls off dramatically at long wavelengths (700 nm).
Copper(II) ions absorb all wavelengths of visible light from 580 to 700 nm. The maximum absorbance occurs at >700 nm, as evidenced by the steeply rising curve. It’s best to choose a wavelength for analysis in the 600–650 nm range because the spectrophotometer may not give accurate measurements at the 700 nm cutoff for visible light.
- If Cu2+ ions and Fe3+ ions are both suspected of being present in the solution, will Fe3+ ions interfere with the analysis of Cu2+ at the wavelength selected? Why or why not? Revise your answer to Question 3, if needed.
Iron(III) ions absorb visible light below 550 nm. Choosing a wavelength above this value for analyzing the concentration of Cu2+ ions in solution will not show any interference due to Fe3+ ions.
- In order to accommodate the range of possible [Cu2+] concentrations that may be obtained by dissolving brass, it’s recommended that the calibration curve extend from approximately 0.05 M to 0.4 M. Estimate the absorbance of standard solutions containing 0.05 M, 0.1 M, 0.2 M and 0.4 M copper(II) nitrate at the selected wavelength. Revise your answer to Question 3, if needed.
We selected 630 nm as the optimum wavelength for analysis of copper ion solutions. The absorbance of the 0.1 M Cu(NO3)2 solution was approximately 0.20 at this wavelength. According to Beer’s law, absorbance is directly proportional to concentration. The estimated absorbance for various Cu2+ standard solutions is shown.
- Calculate the volumes of 0.4 M Cu(NO3)2 stock solution and water required to prepare 8.0 mL of each standard solution for your calibration curve.
The results of dilution calculations to prepare 8.0 mL of each standard solution are included in the table.
- Quality control samples were taken of a batch of nickel–iron alloy produced by Ironic Steel, Inc. in the furnace at their Springfield plant. The alloy must contain 43% nickel, ±0.5%, with the remaining percent iron.
A 1.200 g sample was dissolved in hydrochloric acid and diluted to 100 mL in a volumetric flask. The Beer’s law plot for the absorbance of Fe(NO3)3(aq) versus its concentration is listed.
Calculate the percent iron contained in the alloy sample. Based on your results, is the batch of 43% nickel–iron alloy acceptable?
Mass of alloy sample = 1.200 g
Volume of dissolved alloy solution = 100 mL
Sample absorbance = 0.3730
Slope of Fe3+ calibration curve = 3.4919 M–1
- Iron concentration in alloy solution = (Sample absorbance)/(Slope of Fe3+ calibration curve)
= 0.3730/3.4919 M–1
= 0.107 M
- Moles iron in alloy sample = (Volume of dissolved alloy sample)(Iron concentration of dissolved alloy sample)
= (0.107 M)(0.1 L) = 0.0107 moles
- Percent iron in alloy sample = [(moles Fe)(Molar mass Fe)/grams sample)] x 100
= [(0.0107 moles)(55.85 g/mol)/1.200 g] x 100
Therefore, the amount of nickel is (100 – 49.8)% or 50.2% Ni; the batch is not acceptable.
- The characteristic flame test colors of metal ions are due to atomic emission spectra. Discuss the relationship between the absorption and emission of light and the factors responsible for flame test colors. Include quantization of electron energy levels and Planck’s law in your answer.
When a substance is heated in a flame, the atoms absorb energy from the flame. This absorbed energy allows the electrons to be promoted to excited energy levels. From these excited energy levels, there is a natural tendency for the electrons to make a transition or drop back down to the ground state. When an electron makes a transition from a higher energy level to a lower energy level, a particle of light called a photon is emitted. Both the absorption and emission of energy are quantized—only certain energy levels are allowed.
The energy of each emitted photon is equal to the difference in energy between the excited state and the state to which the electron relaxes. The energy of the emitted photon determines the color of light observed in the flame. The flame color may be described in terms of its wavelength, and Planck’s law may be used to calculate the energy of the emitted photon.
ΔE is the difference in energy between the two energy levels in joules (J).
h is Planck’s constant (h = 6.626 x 10–34 J•sec).
c is the speed of light (c = 2.998 x 108 m/sec).
λ (lambda) is the wavelength of light in meters.
- The wavelength of the characteristic, bright yellow-orange flame test color of sodium is 590 nm. Calculate the average energy (ΔE) associated with this atomic emission line.
Use Planck’s law to calculate the energy difference associated with the atomic emission D line of sodium (590 nm).
Planck’s law: E = hc/λ
h = Planck’s constant, 6.626 x 10−34 J•sec
c = speed of light, 2.998 x 108 m/sec
λ = wavelength in m
||Water, Distilled, 4 L
||Water, Deionized, 4 L
||Flinn Scientific Electronic Balance, 120 x 0.001-g
||Beakers, Polymethylpentene (PMP), 50 mL
||Watch Glass, 100 mm, Borosilicate Glass
||Selected Test Tubes, 13 mm x 100 mm, Box of 12
||Cylinder, Polymethylpentene, 50 mL
||Kimwipes, 15" x 16⅞", Box of 140
||Flinn Multi-Sample Spectrophotometer
||Flask, Volumetric, Borosilicate Glass, 100 mL
||Bottles, Washing, Polyethylene, 500-mL
||Filtering Funnel, Polymethylpentene, 70 mm, Stem Length 150 mm