Introduction to Chemistry Exam Solutions - 3326 Verified Questions

Introduction to Chemistry

Exam Solutions

Course Introduction

Introduction to Chemistry provides a foundational overview of the principles and concepts that form the basis of chemical science. Students will explore the structure of atoms and molecules, properties of matter, the periodic table, chemical reactions, stoichiometry, solutions, and basic thermodynamics. The course emphasizes problem-solving, critical thinking, and laboratory techniques to help students understand the role of chemistry in everyday life and its significance across scientific disciplines. This course is ideal for those beginning their studies in science or preparing for more advanced coursework in chemistry or related fields.

Recommended Textbook

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

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25 Chapters

3326 Verified Questions

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Chapter 1: Units of Measurement for Physical and Chemical Change

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Sample Questions

Q1) A person weighs 77.1 kg. What is his weight in pounds?

A)154 pounds

B)170 pounds

C)35.0 pounds

D)162 pounds

Answer: B

Q2) Ethanol has a boiling point of 351 K. What is the boiling point in °C?

A)78 °C

B)-195 °C

C)122 °C

D)62 °C

E)86 °C

Answer: A

Q3) Gasoline is an example of A)a compound.

B)an element.

C)a heterogeneous mixture.

D)a homogeneous mixture.

Answer: D

Page 3

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Chapter 2: Atoms and Elements

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Sample Questions

Q1) An ion has 8 protons, 9 neutrons, and 10 electrons. The symbol for the ion is

A)17_O2- B)17_O2+ C)19_F+ D)19_F- E)17_Ne2+

Answer: A

Q2) Identify the symbol for silver.

A)S

B)Si

C)Ar

D)Ag

E)Sl

Answer: D

Q3) The atomic number is equal to the number of ________.

Answer: protons

Q4) What group of elements in the periodic table are the most unreactive and why?

Answer: The noble gases are the most unreactive because they do not combine with other elements to form compounds.

Page 4

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Chapter 3: Molecules, Compounds, and Nomenclature

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Sample Questions

Q1) Which of the following has the smallest mass?

A)3.50 × 10²³ molecules of I₂

B)85.0 g of Cl₂

C)2.50 mol of F₂

D)0.050 kg of Br₂

Answer: D

Q2) Which of the following is the correct chemical formula for a molecule of astatine?

A)At

B)At^-

C)At⁺

D)At₂

Answer: D

Q3) Identify a possible molecular formula for C₃H₅ClO.

A)C₆H₁₀ClO₂

B)C₅H₁₀Cl₂O₂

C)C₆H₁₀Cl₂O₂

D)C₆H₁₀O₂

E)C₆H₁₂Cl₂O₂

Answer: C

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Page 5

Chapter 4: Chemical Reactions and Stoichiometry

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Sample Questions

Q1) What is the concentration of NO₃^- ions in a solution prepared by dissolving 25.0 g of Ca(NO₃)₂ in enough water to produce 300. mL of solution?

A)0.254 mol L^-1

B)0.508 mol L^-1

C)0.672 mol L^-1

D)1.02 mol L^-1

Q2) Identify the oxidation state of H in HCl(aq). Mg(s)+ 2HCl(aq) MgCl₂(aq)+ H₂(g)

Q3) What is the concentration (M)of a NaCl solution prepared by dissolving 7.2 g of NaCl in sufficient water to give 425 mL of solution?

Q4) Define a spectator ion.

Q5) What is the concentration (M)of sodium ions in 4.57 L of a .398 mol L^-1 Na₃P solution?

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Chapter 5: Gases

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Sample Questions

Q1) Determine the volume of SO₂ (at STP)formed from the reaction of 96.7 g of FeS₂ and 55.0 L of O₂ (at 398 K and 1.20 bar). The molar mass of FeS₂ is 119.99 g mol^-1. 4FeS₂(s)+ 11O₂(g) 2Fe₂O₃(s)+ 8SO₂(g)

A)36.1 L

B)45.3 L

C)18.1 L

D)27.6 L

E)36.6 L

Q2) How many moles of molecular oxygen are required to produce a pressure of 0.413 bar in a 650 mL container with a temperature of 245 K?

A)0.0132

B)0.199

C)0.00872

D)0.00971

E)0.0245

Q3) Define pressure.

Q4) Define effusion.

Q5) Why does hot air rise?

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Chapter 6: Thermochemistry

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Sample Questions

Q1) Identify the units of heat capacity.

A)J °C^-1

B)J g^-1 °C^-1

C)J mol^-1 °C^-1

D)g °C^-1

E)mol °C^-1

Q2) A piece of uranium with a mass of 12.1 g and an initial temperature of 60.2 °C was placed inside a coffee cup calorimeter. The temperature of the water inside the calorimeter increased from 22.8 °C to 23.9 °C. The specific heat capacity of water is 4.184 J g^-1 °C^-1 and the specific heat capacity of uranium is 0.116 J g^-1 °C^-1. Calculate the mass of water contained inside the calorimeter.

A)1.2 g

B)26 g

C)8.4 g

D)17 g

E)11 g

Q3) Where does the energy absorbed during an endothermic reaction go?

Q4) Explain the difference between H and U.

Q5) 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) For a hydrogen atom, which electronic transition would result in the emission of a photon with the highest energy?

A)2s 3p

B)2p 6d

C)6p 4s

D)7f 5d

Q2) n = 1 to n = 2

A)657 nm

B)7460 nm

C)103 nm

D)1280 nm

E)122 nm

Q3) No two electrons can have the same four quantum numbers. What is this known as?

A)Pauli exclusion principle

B)Hund's rule

C)aufbau principle

D)Heisenberg uncertainty principle

Q4) Define paramagnetic.

Q5) Why don't we observe the wavelength of everyday macroscopic objects?

Q6) Describe how a neon light works.

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Chapter 8: Periodic Properties of the Elements

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Sample Questions

Q1) Which reaction below represents the second electron affinity of S?

A)S(g)+ e S (g)

B)S (g)+ e S² (g)

C)S(g) S (g)+ e

D)S (g) S(g)+ e

E)S² (g) S (g)+ e

Q2) Describe the reaction of the alkali metals with nonmetals.

A)inert

B)vigorous

C)mild reaction

D)forms water

E)dissolves

Q3) An element that has the valence electron configuration 3s¹ belongs to which period and group?

A)period 4; group 7

B)period 4; group 2

C)period 4; group 1

D)period 3; group 2

E)period 3; group 1

Q4) Define electron affinity.

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Chapter 9: Chemical Bonding I: Lewis Theory

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Sample Questions

Q1) Place the following in order of increasing magnitude of lattice energy. MgO LiI CaS

A)CaS < MgO < LiI

B)LiI < CaS < MgO

C)MgO < CaS < LiI

D)LiI < MgO < CaS

E)MgO < LiI < CaS

Q2) Choose the compound below that should have the highest melting point according to the ionic bonding model.

A)SrI₂

B)MgF₂

C)CaCl₂

D)SrF₂

E)SrBr₂

Q3) Which of the following elements can form hypercoordinate compounds?

A)N

B)Br

C)F

D)Be

E)None of the above can form compounds with an expanded octet.

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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) Give the approximate bond angle for a molecule with octahedral electron geometry and square planar molecular geometry.

A)90°

B)180°

C)>120°

D)<90°

E)<120°

Q2) Give the electron geometry, molecular geometry, and hybridization for both carbons in CH₃COOH.

Q3) Identify the number of electron groups around a molecule with a trigonal bipyramidal shape.

A)1

B)2

C)3

D)4

E)5

Q4) Is it possible for a molecule to be nonpolar even though it contains polar bonds? Explain your answer and give an example.

Q5) 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) Define volatile.

Q2) Which of the following is considered a nonbonding atomic solid?

A)Ne

B)Fe

C)I₂

D)Ca

E)Li

Q3) Place the following substances in order of increasing boiling point. Ne Cl₂O₂

A)Ne < Cl₂< O₂

B)Cl₂ < O₂< Ne

C)O₂ < Cl₂< Ne

D)Cl₂ < Ne < O₂

E)Ne < O₂< Cl₂

Q4) Which of the following is considered a molecular solid?

A)Cu

B)NH₄NO₃

C)I₂

D)Xe

E)None of these is a molecular solid.

Page 13

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Chapter 12: Solutions

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Sample Questions

Q1) Calculate the molality of a solution formed by dissolving 27.8 g of LiI in 500.0 mL of water.

A)0.254 mol kg^-1

B)0.394 mol kg^-1

C)0.556 mol kg^-1

D)0.241 mol kg^-1

E)0.415 mol kg^-1

Q2) What volume of a 0.716 mol L^-1 KBr solution is needed to provide 30.5 g of KBr?

A)21.8 mL

B)42.7 mL

C)184 mL

D)357 mL

Q3) How many grams of KBr are required to make 350. mL of a 0.115 mol L^-1 KBr solution?

A)0.338 g

B)3.04 g

C)4.79 g

D)40.3 g

Q4) Define the Tyndall effect.

Page 14

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Chapter 13: Chemical Kinetics

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Sample Questions

Q1) If the concentration of a reactant is 6.25%, how many half-lives has it gone through?

A)7

B)6

C)3

D)4

E)5

Q2) Identify the methods used to monitor a reaction as it occurs in the reaction flask.

A)polarimeter

B)spectrometer

C)pressure measurement

D)none of the above

E)all of the above

Q3) For a reaction, what generally happens if the temperature is increased?

A)a decrease in k occurs, which results in a faster rate

B)a decrease in k occurs, which results in a slower rate

C)an increase in k occurs, which results in a faster rate

D)an increase in k occurs, which results in a slower rate

E)there is no change to either k or the rate

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Chapter 14: Chemical Equilibrium

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Sample Questions

Q1) Consider the following reaction: CuS(s)+ O₂(g) Cu(s)+ SO₂(g)

A reaction mixture initially contains 2.9 mol L^-1 O₂. Determine the equilibrium concentration of O₂ if K_c for the reaction at this temperature is 1.5.

A)1.9 mol L^-1

B)1.7 mol L^-1

C)2.2 mol L^-1

D)1.2 mol L^-1

E)0.59 mol L^-1

Q2) In a reaction mixture containing reactants and products, each at a concentration of 1 mol L^-1, what is the value of Q?

A)-1

B)1

C)

D)0

E)Q = K

Q3) How will equilibrium be affected for a reaction having the same number of moles of gaseous reactants and products if the pressure is increased?

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Chapter 15: Acids and Bases

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Sample Questions

Q1) Which of the following acids is the strongest? The acid is followed by its K_a value.

A)HF, 3.5 × 10^-4

B)HCN, 4.9 × 10^-10

C)HNO₂, 4.6 × 10^-4

D)HCHO₂, 1.8 × 10^-4

E)HClO₂, 1.1 × 10^-2

Q2) When dissolved in water, which compound is generally considered to be an Arrhenius acid?

A)CH₃CO₂H

B)NaOH

C)Na₂CO₃

D)CH₃CH₂OH

Q3) Calculate the pH of a 0.60 mol L^-1 H₂SO₃ solution that has the stepwise dissociation constants K_a1 = 1.5 × 10^-2 and K_a2 = 6.3 × 10^-8.

A)1.02

B)1.06

C)1.82

D)2.04

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Chapter 16: Aqueous Ionic Equilibrium

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Sample Questions

Q1) Identify the indicator that has two endpoints.

A)phenol red

B)thymol blue

C)crystal violet

D)phenolphthalein

E)alizarin yellow R

Q2) A solution containing AgNO₃ is mixed with a solution of NaCl to form a solution that is 0.10 mol L^-1 in AgNO₃ and 0.075 mol L^-1 in NaCl. What will happen once these solutions are mixed?

K_sp (AgCl)= 1.77 × 10^-10.

A)Nothing will happen since the molar solubility of AgCl is higher than the solution concentrations.

B)Silver chloride will precipitate out of the solution, leaving an unsaturated solution of AgCl.

C)Silver chloride will precipitate out of the solution, leaving a saturated AgCl solution.

D)Nothing will happen since NaCl and AgNO₃ are both soluble compounds.

E)Silver chloride will precipitate, leaving a pure solution of NaNO₃.

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Chapter 17: Gibbs Energy and Thermodynamics

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Sample Questions

Q1) Which of the following processes have a _rS > 0?

A)CH₃OH(l) CH₃OH(s)

B)N₂(g)+ 3H₂(g) 2NH₃(g)

C)CH₄(g)+ H₂O (g) CO(g)+ 3H₂(g)

D)Na₂CO₃(s)+ H₂O(g)+ CO₂(g)

2NaHCO₃(s)

E)H₂O(g) H₂O(l)

Q2) The value of _fG° at 100.0 °C for the formation of calcium chloride from its constituent elements, Ca(s)+ Cl₂(g) CaCl₂(s)

Is __________ kJ mol^-1. At 25.0 °C for this reaction, _fH° is -795.8 kJ mol^-1l, _fG° is -748.1 kJ mol^-1, and _fS° is -159.8 J K^-1 mol^-1.

A)-855.4

B)-736.2

C)5.88 × 10⁴

D)-779.8

E)1.52 × 10⁴

Q3) Give the standard states for a gas, liquid, solid, and solution.

Q4) Why can endothermic reactions be spontaneous?

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Chapter 18: Electrochemistry

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Sample Questions

Q1) Describe how water can be made to be a good conductor of electrical current.

A)use pure water

B)heat the water

C)add salt

D)chill the water

E)vaporize the water

Q2) What is the shorthand notation that represents the following galvanic cell reaction?

2Fe²⁺(aq)+ F₂(g) 2Fe³⁺(aq)+ 2F^-(aq)

A)Fe²⁺(aq) Fe³⁺(aq) F₂(g) F^-(aq)

B)Fe(s) Fe²⁺(aq) Fe³⁺(aq)F₂(g) F^-(aq) C(s)

C)Pt(s) Fe³⁺(aq), Fe²⁺(aq), F₂(g) F^-(aq) C(s)

D)Pt(s) Fe²⁺(aq), Fe³⁺(aq) F₂(g) F^-(aq) C(s)

Q3) Explain the significance of the standard hydrogen electrode (SHE)in the tabulation of standard reduction potentials of other species.

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

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Chapter 19: Radioactivity and Nuclear Chemistry

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Sample Questions

Q1) Calculate the mass defect in Mo-96 if the mass of a Mo-96 nucleus is 95.962 amu. The mass of a proton is 1.00728 amu and the mass of a neutron is 1.008665 amu.

A)0.197 amu

B)0.795 amu

C)0.212 amu

D)0.812 amu

E)0.188 amu

Q2) Strontium-90 is a byproduct in nuclear reactors fuelled by the radioisotope uranium-235. The half-life of strontium-90 is 28.8 y. What percentage of a strontium-90 sample remains after 75.0 y.?

A)68.1

B)16.4

C)7.40

D)38.4

E)2.60

Q3) Explain the concept of "magic numbers."

Q4) Define radioactivity.

Q5) Describe what is meant by the "valley of stability."

Q6) How does a dosimeter measure exposure to radioactivity?

Page 21

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Chapter 20: Organic Chemistry I: Structures

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Sample Questions

Q1) What is the difference between primary, secondary, and tertiary amines?

Q2) Why are there so many more carbon compounds than the number of compounds made up of all the rest of the elements combined?

Q3) Which one of the following molecules is the most polar?

A)butane

B)acetic acid

C)cyclohexane

D)methanol

E)cyclohexanone

Q4) Which of the following condensed general formulas represents ethers?

A)ROR

B)RCOR

C)RCHO

D)ROH

E)RCOOH

Q5) What is meant by the term, "structural isomer"? Draw a structural isomer of CH₃-CH₂-CH₂-CH₃.

Q6) What is the difference between the R/ S designation and d/l designation of chiral compounds?

Page 22

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Chapter 21: Organic Chemistry II: Reactions

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Sample Questions

Q1) Which of the following statements is TRUE?

A)Addition is the predominant reaction observed in acyl compounds.

B)Acyl chlorides are the least reactive of all acyl compounds, while amides are the most.

C)Fischer esterification is the reaction between carboxylic acids with alcohols to produce esters.

D)Fischer esterification is catalyzed by a strong base.

E)Transesterification is a reaction in which an ester group is transferred from one organic group to another.

Q2) Give the organic product for the following reaction.

CH₃CH₂CH₂NH₂ + HCl

A)ClCH₂CH₂CH₂NH₂ B)CH₃CHClCH₂NH₂

C)CH₃CH₂CHClNH₂ D)CH₃CH₂CH₂NHCl

E)CH₃CH₂CH₂NH₃⁺Cl<su p>-</sup>

Q3) Why doesn't benzene typically undergo addition reactions like the alkenes do?

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Chapter 22: Biochemistry

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Sample Questions

Q1) Identify the disaccharide.

A)glucose

B)cholesterol

C)fructose

D)cellulose

E)sucrose

Q2) Why are lipids well-suited as structural components of cell membranes?

Q3) The following is part of a DNA sequence. What is its complementary sequence?

TCGTAAGCTTGCG

A)AGCATTCGAACGC

B)CGCAAGCTTACGA

C)TCATTCAC

D)GCGTTCGAATGCT

E)TCGTAAGCTTGCG

Q4) Which of the following link together into amino acid units?

A)glycosidic linkages

B)disulfide linkages

C)peptide bonds

D)hydrogen bonds

E)ester linkages

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Chapter 23: Chemistry of the Nonmetals

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Sample Questions

Q1) Determine the number of vertices and faces in the closo-Borane B₁₀H₁₀^2-.

A)vertices = 10, faces = 16

B)vertices = 5, faces = 6

C)vertices = 5, faces = 10

D)vertices = 10, faces = 20

E)vertices = 20, faces = 36

Q2) Which of the following statements is TRUE?

A)Boron is a large component of Earth's crust.

B)Boron is usually found in its elemental state in the Earth's crust.

C)Due to boron's small size and low electronegativity, it behaves as a semimetal instead of a metal.

D)Boron atoms don't typically bond to one another due to its small atomic radius.

E)There are at least five allotropes of elemental boron.

Q3) Why are the chemical properties of nitrogen and phosphorus so different when they are in the same family?

Q4) Describe the major production method for obtaining oxygen.

Q5) Why are phosphate compounds added to detergents?

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Chapter 24: Metals and Metallury

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Q1) Smelting is

A)heating an ore in the presence of oxygen or another substance to purify the ore and obtain the liquid metal.

B)selectively dissolving a metal in solution to separate it from its ore.

C)the process of heating an ore to drive off volatile compounds.

D)heating an ore in the presence of oxygen or another substance to cause a chemical reaction that drives off newly formed volatile compounds.

E)forming metal parts using heat and small crystals of metal.

Q2) Hg

A)sphalerite

B)rhodochrosite

C)rutile

D)carnotite

E)malachite

F)galena

G)cinnabar

Q3) Why does a "two-phase" structure occur in a substitutional alloy?

Q4) Why is zinc used to coat steel objects?

Q5) What is the difference between ferromagnetism and paramagnetism?

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Chapter 25: Transition Metals and Coordination Compounds

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Sample Questions

Q1) Which of the following can function as a bidentate ligand?

A)CN^-

B)SH^-

C)HO

D)H₂NCH₂

Q2) How many unpaired electrons would you expect for the complex ion

[Fe(CN)₆]^4-?

A)5

B)2

C)6

D)4

E)0

Q3) Which ion would you expect to have the largest crystal field splitting, ?

A)[Rh(CN)₆]^4-

B)[Rh(CN)₆]^3-

C)[Rh(H₂O)₆]²⁺

D)[Rh(H₂O)₆]³⁺

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

Page 27

Q5) What is a coordinate covalent bond?

Q6) What is the difference between a weak-field complex and a strong-field complex?

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