Chemistry for Engineers Study Guide Questions - 3326 Verified Questions

Course Introduction

Chemistry for Engineers Study Guide

Questions

This course introduces the fundamental principles of chemistry with an emphasis on concepts and applications relevant to engineering. Topics include atomic and molecular structure, chemical bonding, stoichiometry, thermodynamics, kinetics, equilibrium, and electrochemistry. Students will explore how chemical processes and properties influence material selection, technological advancements, and environmental considerations in various engineering fields. Laboratory sessions and problem-solving exercises are integrated to develop practical skills and analytical reasoning necessary for addressing real-world challenges faced by engineers.

Recommended Textbook

Chemistry A Molecular Approach 1st Canadian Ediiton by Nivaldo J. Tro

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Page 2

Chapter 1: Units of Measurement for Physical and Chemical Change

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Sample Questions

Q1) Which of the following statements about the phases of matter is TRUE?

A)In both solids and liquids, the atoms or molecules pack closely to one another.

B)Solids are highly compressible.

C)Gaseous substances have long-range repeating order.

D)There is only one type of geometric arrangement that the atoms or molecules in any solid can adopt.

E)Liquids have a large portion of empty volume between molecules.

Answer: A

Q2) Convert 4 m to metres.

A)4 × 10^-9 m

B)4 × 10^-6 m

C)4 × 10^-3 m

D)4 × 10⁶ m

Answer: B

Q3) A flash drive contains 4 gigabytes. How many bytes does it contain?

Answer: 4 000 000 000 bytes, or 4 292 967 296 bytes if you are is computer literate

Q4) 38.325 lbs = ________ grams. (1 lb = 454 g)

Answer: 17400

Page 3

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Chapter 2: Atoms and Elements

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Sample Questions

Q1) Which of the following describes a metal?

A)Poor conductor of heat

B)Good conductor of electricity

C)Tends to gain electrons

D)Forms ionic compounds with group 18 elements

E)Found on the upper right corner of the periodic table

Answer: B

Q2) How many atoms are in 1.00 kg of copper?

A)3.83 × 10²⁹ atoms

B)3.83 × 10²² atoms

C)15.74 atoms

D)2.61 × 10^-23 atoms

E)9.48 × 10²⁴ atoms

Answer: E

Q3) Give the name of the element whose symbol is Na.

Answer: sodium

Q4) Why do elements in the same group tend to have similar chemical properties?

Answer: Since elements in the same group have the same number of valence electrons (similar electron configurations)they tend to have similar chemical reactivity because chemical reactions typically involve valence electrons.

Page 4

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Chapter 3: Molecules, Compounds, and Nomenclature

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Sample Questions

Q1) Methane and oxygen react to form carbon dioxide and water. What mass of water is formed if 0.80 g of methane reacts with 3.2 g of oxygen to produce 2.2 g of carbon dioxide?

A)1.8 g

B)2.2 g

C)3.7 g

D)4.0 g

Answer: A

Q2) Calculate the molar mass of Ca₃(PO₄)₂.

A)87.05 g mol^-1

B)215.21 g mol^-1

C)310.18 g mol^-1

D)279.21 g mol^-1

E)246.18 g mol^-1

Answer: C

Q3) Determine the mass percent (to the hundredths place)of H in sodium bicarbonate (NaHCO₃).

Answer: 1.20

Q4) Define empirical formula.

Answer: An empirical formula gives relative numbers of atoms of each element.

Page 5

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Chapter 4: Chemical Reactions and Stoichiometry

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Sample Questions

Q1) Determine the concentration of a solution prepared by diluting 20.0 mL of 0.200 mol L^-1 NaCl to 250.0 mL.

A)0.160 mol L^-1

B)0.0320 mol L^-1

C)2.50 mol L^-1

D)0.00800 mol L^-1

E)0.0160 mol L^-1

Q2) What is the stoichiometric coefficient for oxygen when the following equation is balanced using the lowest, whole-number coefficients? _____ CH₄O(l)+ _____ O₂(g) _____ CO₂(g)+ _____ H₂O(l)

A)9

B)7

C)5

D)3

Q3) Which of the following compounds is an Arrhenius base?

A)C₆H₁₂O₆

B)HOCl

C)H₂SO₄

D)C₆H₅NH₂

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Chapter 5: Gases

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Sample Questions

Q1) Determine the density of O₃ gas at 341 K and 2.14 bar.

A)3.62 g L^-1

B)2.91 g L^-1

C)0.321 g L^-1

D)4.82 g L^-1

E)3.17 g L^-1

Q2) What pressure (in bar)will 0.44 moles of CO₂ exert in a 2.6 L container at 25 °C?

A)0.35 bar

B)4.2 bar

C)4.7 bar

D)8.6 bar

E)3.6 bar

Q3) How many moles of CO are contained in a 5.00 L tank at 155 °C and 2.80 bar?

A)0.393 moles

B)1.10 moles

C)2.51 moles

D)0.455 moles

E)0.289 moles

Q4) Define pressure.

Page 7

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Chapter 6: Thermochemistry

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Sample Questions

Q1) Identify an energy source that is NOT renewable.

A)solar

B)hydroelectric

C)coal

D)wind

E)sun

Q2) Choose the thermochemical equation that illustrates _fH° for Li₂SO₄.

A)2Li⁺(aq)+

Li₂SO₄(aq)

SO₄^2-(aq)

B)2Li(s)+ 1/8 S₈(s, rhombic)+

2O₂(g) Li₂SO₄(s)

C)Li₂SO₄(aq)

2Li⁺(aq)+ SO₄^2-(aq)

D)8Li₂SO₄(s) 16Li(s)+ S₈(s, rhombic)+ 16O₂(g)

E)16Li(s)+ S₈(s, rhombic)+ 16O₂(g) 8Li₂SO₄(s)

Q3) Define chemical energy.

Q4) Give the temperature and pressure for the standard state for a liquid.

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Chapter 7: The Quantum-Mechanical Model of the Atom

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Sample Questions

Q1) Only two electrons, with opposing spins, are allowed in each orbital. What is this known as?

A)Pauli exclusion principle

B)Hund's rule

C)aufbau principle

D)Heisenberg uncertainty principle

Q2) What is the maximum number of p orbitals that are possible?

A)1

B)3

C)7

D)5

E)9

Q3) Define constructive interference.

Q4) How many of the following species are paramagnetic? Sc³ Br Mg² Se

A)0

B)2

C)1

D)4

E)3

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Chapter 8: Periodic Properties of the Elements

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Sample Questions

Q1) List the noble gas that has the highest ionization energy.

Q2) Using the periodic table, identify the element with the following exception electron configuration: [Ar]4s¹3d⁵

A)Ni

B)Co

C)Cr

D)Mn

E)Fe

Q3) Which ionization process requires the most energy?

A)W(g) W⁺(g)+ e^-

B)W⁺(g) W²⁺(g)+ e^-

C)W²⁺(g) W³⁺(g)+ e^-

D)W³⁺(g) W⁴⁺(g)+ e^-

Q4) Which element has the highest first electron affinity?

A)Na

B)Mg

C)O

D)Ne

Q5) Why does the size of the transition elements stay roughly the same as you move across a period?

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Chapter 9: Chemical Bonding I: Lewis Theory

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Sample Questions

Q1) Use the bond energies provided to estimate _rH° for the reaction below.

XeF_{2 +}2F₂ XeF₆ _rH°= ?

Bond Bond Energy (kJ mol^-1)

Xe-F 147

F-F 159

A)-429 kJ mol^-1

B)+159 kJ mol^-1

C)-660 kJ mol^-1

D)+176 kJ mol^-1

E)-270 kJ mol^-1

Q2) Identify the compound with covalent bonding.

A)NaCl

B)Li

C)H₂O

D)He

E)S

Q3) Define formal charge.

Q4) List the most electronegative atom.

Q5) How are electron affinity and electronegativity different?

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Chapter 10: Chemical Bonding Ii: Molecular Shapes,

Valence Bond Theory, and Molecular Orbital Theory

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Sample Questions

Q1) How many of the following molecules are polar?

XeO₂ SiCl₂Br₂ C₂Br₂ SeCl₆

A)1

B)4

C)2

D)3

E)0

Q2) What is the molecular geometry of NCl₃?

A)T-shaped

B)tetrahedral

C)trigonal planar

D)trigonal pyramidal

Q3) Draw the Lewis structure for BrF₅. What is the hybridization on the Br atom?

A)sp³d²

B)sp³d

C)sp³

D)sp²

E)sp

Q4) Explain why oil and water do not mix.

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Chapter 11: Liquids, Solids, and Intermolecular Forces

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Sample Questions

Q1) Describe the difference between the conduction band and the valence band.

Q2) Why is water an extraordinary substance?

A)Water has a low molar mass, yet it is a liquid at room temperature.

B)Water is the main solvent within living organisms.

C)Water has an exceptionally high specific heat capacity.

D)Water has strong hydrogen bonding.

E)All of the above.

Q3) Which of the following substances should have the highest melting point?

A)CO₂

B)SrS

C)Xe

D)F₂

E)MgO

Q4) Identify the type of solid for argon.

A)metallic atomic solid

B)ionic solid

C)nonbonding atomic solid

D)molecular solid

E)networking atomic solid

Q5) Define viscosity.

Page 13

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Chapter 12: Solutions

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Sample Questions

Q1) Explain why the van't Hoff factor for MgCl₂ is less than its predicted value.

Q2) Calculate the total solution volume required to produce an osmotic pressure of 4.87 mbar using 8.21 mg of a protein with a molar mass of 4450 g mol^-1 at 25.0 °C.

A)10.6 mL

B)4.79 mL

C)12.3 mL

D)9.39 mL

E)8.00 mL

Q3) What is the mole fraction of ethanol in a solution made by dissolving 29.2 g of ethanol, C₂H₅OH, in 53.6 g of water?

A)0.176

B)0.213

C)0.352

D)0.545

Q4) Explain why water does not dissolve in gasoline.

Q5) Give the preparation of rock candy.

Q6) Define the Tyndall effect.

Q7) Define a colloid.

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Chapter 13: Chemical Kinetics

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Sample Questions

Q1) The first-order decay of radon has a half-life of 3.823 days. How many grams of radon decomposes after 5.55 days if the sample initially weighs 100.0 grams?

A)83.4 g

B)16.6 g

C)50.0 g

D)36.6 g

E)63.4 g

Q2) The combustion of ethylene proceeds by the reaction C₂H₄(g)+ 3O₂(g) 2CO₂(g)+ 2H₂O(g)

When the rate of disappearance of O₂ is 0.23 mol L^-1 s^-1, the rate of disappearance of C₂H₄ is ________ mol L^-1 s^-1.

A)0.15

B)0.077

C)0.69

D)0.35

E)0.46

Q3) Define the frequency factor.

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Page 15

Chapter 14: Chemical Equilibrium

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Sample Questions

Q1) Consider the following reaction: 2H₂O(g)+ 2SO₂(g) 2H₂S(g)+ 3O₂(g)

A reaction mixture initially contains 2.8 mol L^-1 H₂O and 2.6 mol L^-1 SO₂. Determine the equilibrium concentration of H₂S if K_c for the reaction at this temperature is 1.3 × 10^-6mol L^-1.

A)0.045 mol L^-1

B)0.058 mol L^-1

C)0.028 mol L^-1

D)3.1 × 10^-3 mol L^-1

E)0.12 mol L^-1

Q2) For the following reaction, what is n required in the conversion of K_c to K_p? N₂O₄(g) 2NO₂(g)

Q3) Why is a thermodynamic equilibrium constant unitless?

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Chapter 15: Acids and Bases

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Sample Questions

Q1) Which of the following is an Arrhenius base?

A)CH₃CO₂H

B)NaOH

C)CH₃OH

D)LiCl

E)H₂CO₃

Q2) Which of the following is the acid found in our stomachs?

A)H₂SO₄

B)HI

C)HCl

D)HF

E)HBr

Q3) Determine the pH of a 0.18 mol L^-1 H₂CO₃ solution. Carbonic acid is a diprotic acid whose K_a1 = 4.3 × 10^-7 and K_a2 = 5.6 × 10^-11.

A)11.00

B)10.44

C)5.50

D)4.31

E)3.56

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Chapter 16: Aqueous Ionic Equilibrium

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Sample Questions

Q1) When titrating a weak monoprotic acid with NaOH at 25 °C, the A)pH will be less than 7 at the equivalence point.

B)pH will be equal to 7 at the equivalence point.

C)pH will be greater than 7 at the equivalence point.

D)titration will require more moles of base than acid to reach the equivalence point.

E)titration will require more moles of acid than base to reach the equivalence point.

Q2) What is the pH of a solution prepared by mixing 25.00 mL of 0.10 mol L^-1 CH₃COOH with 25.00 mL of 0.010 mol L^-1 CH₃COONa? Assume that the volume of the solutions are additive and that K_a = 1.8 × 10^-5 for CH₃COOH.

A)2.87

B)3.74

C)4.75

D)5.74

Q3) Explain the common ion effect with respect to molar solubility.

Q4) Define a buffer.

Q5) Give the name of the compound that is in antifreeze and is toxic to pets.

Q6) Identify the compound that forms stalactites and stalagmites.

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Chapter 17: Gibbs Energy and Thermodynamics

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Sample Questions

Q1) Calculate S° for the following reaction. The S° for each species is shown below the reaction. P₄(g)+ 10Cl₂(g) 4PCl₅(g)

S°(J K^-1 mol^-1)280.0 223.1 364.6

A)-138.5 J K^-1 mol^-1

B)-1052.6 J K^-1 mol^-1

C)+171.3 J K^-1 mol^-1

D)-583.6 J K^-1 mol^-1

E)+2334.6 J K^-1 mol^-1

Q2) Calculate _rG° at 298 K using the following information: 4HNO₃(g)+ 5N₂H₄(l) 7N₂(g)+ 12H₂O(l) _rG°= ? _fG° (kJ mol^-1)-73.5 149.3 -237.1

A)-3.2977 × 10³ kJ mol^-1

B)-312.9 kJ mol^-1

C)+2.845 × 10³ kJ mol^-1

D)+110.7 kJ mol^-1

E)-954.7 kJ mol^-1

Q3) Define the third law of thermodynamics.

Q4) What is "free" energy? Give a fictitious example.

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Chapter 18: Electrochemistry

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Sample Questions

Q1) Predict the species that will be reduced first if a mixture of molten salts containing the following ions undergoes electrolysis: Zn²⁺, Fe³⁺, Mg²⁺, Br^-, I^-

A)Zn²⁺

B)Mg²⁺

C)Br^-

D)Fe³⁺

E)I^-

Q2) Based on the following information, Cl₂(g)+ 2 e^-

2Cl^-(aq)E° = +1.36 V

Mg²⁺(aq)+ 2 e^- 2Mg(s)E° = -2.37 V

Which of the following chemical species is the strongest reducing agent?

A)Cl₂(g)

B)Mg²⁺(aq)

C)Cl^-(aq)

D)Mg(s)

Q3) Why is sugar water not a good conductor of current?

Q4) What is the difference between a voltaic cell and an electrolytic cell?

Q5) What is electrolysis?

Q6) Give an example of an inert electrode.

Page 20

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Chapter 19: Radioactivity and Nuclear Chemistry

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Sample Questions

Q1) Describe what changes occur in the atomic nucleus during beta decay.

A)The mass number and atomic number decrease.

B)The mass number and atomic number increase.

C)The mass number is unchanged and the atomic number decreases.

D)The mass number is unchanged and the atomic number increases.

E)The mass number and atomic number do not change.

Q2) Calculate the mass defect in Fe-56 if the mass of an Fe-56 nucleus is 55.921 amu.

The mass of a proton is 1.00728 amu and the mass of a neutron is 1.008665 amu.

A)0.528 amu

B)3.507 amu

C)0.564 amu

D)1.056 amu

E)0.079 amu

Q3) Determine the half-life of a nuclide that loses 38.0% of its mass in 387 hours.

A)277 hours

B)455 hour

C)561 hours

D)639 hours

E)748 hours

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Page 21

Chapter 20: Organic Chemistry I: Structures

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Sample Questions

Q1) What is the index of hydrogen deficiency?

Q2) Which of the following compounds is ethanol?

A)C₂H₆

B)C₂H₅OH

C)CH₃CO₂H

D)CH₃CO₂CH₃

E)CH₃-O-CH₃

Q3) Which of the following statements is TRUE?

A)Infrared absorption spectroscopy is a common method for identifying functional groups in the structure of organic molecules.

B)Infrared absorption spectroscopy is a technique that makes use of the fact that infrared photons can excite the valence electrons in organic molecules.

C)Infrared absorption spectroscopy is a simple yet powerful method for determining the absolute configuration of chiral organic molecules.

D)Infrared absorption spectroscopy observes the energy emitted by vibrating bonds in molecules.

E)Infrared absorption spectroscopy is a technique frequently used in separation of isomers.

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Chapter 21: Organic Chemistry II: Reactions

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Q1) Which of the following statements is TRUE?

A)The S_N2 mechanism does not have any intermediates and has only one transition state.

B)Tertiary alkylchlorides are the best substrates for S_N2 reactions.

C)Only the strength of the nucleophile is important for S_N2 reactions, not the nature of the leaving group.

D)S_N1 reactions have a carbanion as an intermediate.

E)S_N2 reactions result in the racemic product if we start from one enantiomer of a chiral substrate

Q2) In which of the following families would the strongest organic bases be found?

A)amines

B)ethers

C)esters

D)alcohols

E)ketones

Q3) Which part of HCN, the hydrogen or the cyano group, adds to the O in a C=O bond and why?

Q4) Why doesn't benzene typically undergo addition reactions like the alkenes do?

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Chapter 22: Biochemistry

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Q1) Which one of the following amino acids does not contain a basic side chain?

A)arginine

B)histidine

C)lysine

D)threonine

Q2) What is the appropriate name for the strongest intermolecular force holding together two strands of DNA to form a double helix?

A)hydrogen bonds

B)London dispersion forces

C)dipole-dipole bonds

D)ionic bonds

E)covalent bonds

Q3) Which of the following is a type of nucleic acid?

A)RNA

B)amino acid

C)dipeptide

D)carbohydrate

E)lipid

Q4) How do cis- and trans-fats differ?

Q5) What is a codon?

24

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Chapter 23: Chemistry of the Nonmetals

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Q1) Tetraphosphorus decoxide reacts with water to produce phosphoric acid. Write a balanced reaction for this process.

A)P₄O₁₀ + 6H₂O 4H₃PO₄ B)P₂O₅ + 3H₂O 2H₃PO₃

C)P₄ + 3O_{2 +}6H₂O 4H₃PO₄

D)P₄O₁₀ + 4H₂O 4H₂PO₃

E)P₄ + 3O_{2 +}6H₂O 4H₃PO₃

Q2) What is the shape of a boron trihalide?

A)trigonal planar

B)tetrahedral

C)linear

D)trigonal pyramidal

E)bent

Q3) Why are phosphate compounds added to detergents?

Q4) Draw the Lewis structure for IF₄ and determine the geometry of the ion.

Q5) Why are the chemical properties of nitrogen and phosphorus so different when they are in the same family?

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Chapter 24: Metals and Metallury

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Q1) Why can't chromium and nickel form a miscible solid solution over the entire composition range?

A)The lever rule says that these two metals can't coexist in an alloy.

B)The two metals crystallize into different cubic structures in their pure forms.

C)The tetrahedral holes in the nickel crystal are not large enough to accommodate the chromium.

D)The octahedral holes in the chromium crystal are not large enough to accommodate the nickel.

E)Chromium atoms can only fit inside tetrahedral holes, which are not present in the structure of pure nickel.

Q2) Which of the following processes is used to form metal components from very small metal particles?

A)electrometallurgy

B)hydrometallurgy

C)pyrometallurgy

D)powder metallurgy

E)smelting

Q3) Why is zinc used to coat steel objects?

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Chapter 25: Transition Metals and Coordination Compounds

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Sample Questions

Q1) Choose the electron configuration for Nb³⁺.

A)[Kr]5s¹4d²

B)[Kr]5s²4d¹

C)[Kr]5s²

D)[Kr]4d²

E)[Kr]5s²4d⁵

Q2) Choose the electron configuration for Ti²⁺.

A)[Ar]3d²

B)[Ar]4s²3d²

C)[Ar]4s²3d⁴

D)[Ar]4s²

E)[Ne]3s²3p⁶

Q3) Which of the following complex ions absorbs light of the longest wavelength?

A)[Cr(CN)₆]^3-

B)[CrCl₆]^3-

C)[Cr(en)₃]³⁺

D)[Cr(NH₃)₆]³⁺

E)[Cr(NO₂)₆]^3-

Q4) Explain how EDTA is used to treat lead poisoning.

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Q5) What is the difference between a weak-field complex and a strong-field complex?