Chemistry for Life Sciences
Midterm Exam
Course Introduction
Chemistry for Life Sciences provides a foundational understanding of chemical principles with a focus on their applications in biological systems. The course introduces students to atomic structure, chemical bonding, thermodynamics, acids and bases, and the properties of biomolecules such as proteins, carbohydrates, lipids, and nucleic acids. Emphasis is placed on how chemical reactions and interactions drive processes critical to life, including metabolism, cellular communication, and genetic information transfer. Through lectures and laboratory work, students develop practical skills and analytical approaches essential for studies in biology, medicine, and related health sciences.
Recommended Textbook
Chemistry A Molecular Approach 2nd Canadian Edition by Nivaldo J. Tro
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Chapter 1: Units of Measurement for Physical and Chemical Change
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Sample Questions
Q1) What answer should be reported,with the correct number of significant figures,for the following calculation?
(249.362 + 42)/ 63.498
A)4.6
B)4.59
C)4.589
D)4.5885
E)4.58852
Answer: B
Q2) A student performs an experiment to determine the density of a sugar solution.She obtains the following results: 1.71 g mL<sup>-1</sup>,1.73 g mL<sup>-1</sup>,1.67 g mL<sup>-1</sup>,1.69 g mL<sup>-1</sup>.If the actual value for the density of the sugar solution is 1.40 g mL<sup>-1</sup>,which statement below best describes her results?
A)Her results are precise,but not accurate.
B)Her results are accurate,but not precise.
C)Her results are both precise and accurate
D)Her results are neither precise nor accurate.
E)It isn't possible to determine from the information given.
Answer: A
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Chapter 2: Atoms and Elements
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Sample Questions
Q1) Which of the following elements is a halogen?
A)Ne
B)I
C)O
D)Mg
E)K
Answer: B
Q2) Dalton's atomic theory states
A)that all elements have several isotopes.
B)that matter is composed of small indestructible particles.
C)that the properties of matter are determined by the properties of atoms.
D)that energy is neither created nor destroyed during a chemical reaction.
E)that an atom is predominantly empty space.
Answer: B
Q3) Give an example of a halogen.
Answer: F,Br,I,Cl,or At
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.
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Chapter 3: Molecules,compounds,and Nomenclature
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Sample Questions
Q1) How many molecules are in 114.86 g of vitamin C,C<sub>6</sub>H<sub>8</sub>O<sub>6</sub>?
A)1.218 × 10<sup>28</sup>
B)3.927 × 10<sup>23</sup>
C)2.546 × 10<sup>-24</sup>
D)2.977 × 10<sup>19</sup>
E)2.977 × 10<sup>22</sup>
Answer: B
Q2) How many moles of C<sub>3</sub>H<sub>8</sub> contain 9.25 × 10<sup>24</sup> molecules of C<sub>3</sub>H<sub>8</sub>?
A)65.1<sub> </sub>moles C<sub>3</sub>H<sub>8</sub>
B)28.6<sub> </sub>moles C<sub>3</sub>H<sub>8</sub>
C)34.9<sub> </sub>moles C<sub>3</sub>H<sub>8</sub>
D)46.2<sub> </sub>moles C<sub>3</sub>H<sub>8</sub>
E)15.4 moles C<sub>3</sub>H<sub>8</sub>
Answer: E
Q3) Give the name for HNO<sub>2</sub>.
Answer: nitrous acid
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) Identify the location of reduction in an electrochemical cell.
A)the anode
B)the cathode
C)the electrode
D)the salt bridge
E)the socket
Q2) When 280.mL of 1.50 × 10<sup>-4 </sup>mol L<sup>-1</sup> hydrochloric acid is added to 135 mL of 1.75 × 10<sup>-4</sup> mol L<sup>-1</sup> Mg(OH)<sub>2</sub>,the resulting solution will be
A)acidic.
B)basic
C)neutral.
D)It is impossible to tell from the information given.
Q3) Which of the following pairs of aqueous solutions will form a precipitate when mixed?
A)K<sub>2</sub>CO<sub>3</sub> + NaCl
B)Na<sub>2</sub>SO<sub>4</sub> + KOH
C)CaS + Na<sub>2</sub>SO<sub>4</sub>
D)None of these solution pairs will produce a precipitate.
E)All of these solution pairs will produce a precipitate.
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Chapter 5: Gases
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Sample Questions
Q1) Determine the total volume of all gases (at STP)formed when 50.0 mL of TNT C<sub>3</sub>H<sub>5</sub>(NO<sub>3</sub>)<sub>3</sub>,d = 1.60 g mL<sup>-1</sup>,molar mass = 227.10 g mol<sup>-1</sup>)reacts according to the following reaction:
4C<sub>3</sub>H<sub>5</sub>(NO<sub>3</sub>)<sub>3</sub>(l) 6N<sub>2</sub>(g)+ O<sub>2</sub>(g)+ 12CO<sub>2</sub>(g)+ 10H<sub>2</sub>O(g)
A)4.93 L
B)58.0 L
C)29.6 L
D)448 L
E)175 L
Q2) Calculate the root mean square velocity of nitrogen molecules at 25 °C.
A)729 m s<sup>-1</sup>
B)515 m s<sup>-1</sup>
C)149 m s<sup>-1</sup>
D)297 m s<sup>-1</sup>
Q3) Define effusion.
Q4) Why does hot air rise?
Q5) Define pressure.
Q6) Why does the rate of effusion increase with a decrease in the molar mass?
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Chapter 6: Thermochemistry
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Sample Questions
Q1) Using the following equation for the combustion of octane,calculate the heat of reaction for 100.0 g of octane.The molar mass of octane is 114.33 g \(\mathrm { mol } ^ { - 1 }\) . 2C<sub>8</sub>H<sub>18</sub> + 25O<sub>2</sub> ? 16CO<sub>2</sub> + 18H<sub>2</sub>O
D<sub>r</sub>H°<sub> </sub>= -11018 kJ
A)4.82 × 10<sup>3</sup> kJ
B)4.82 kJ
C)9.64 × 10<sup>3 </sup>kJ
D)1.26 × 10<sup>4 </sup>kJ
Q2) How much heat is absorbed when 50.00 g of C(s)reacts in the presence of excess SO<sub>2</sub>(g)to produce CS<sub>2</sub>(l)and CO(g)according to the following chemical equation? 5C(s)+ 2SO<sub>2</sub>(g)? CS<sub>2</sub>(l)+ 4CO(g) \(\Delta _ { \mathrm { r } }\) H° = 239.9 kJ
A)199.7 kJ \(\mathrm { mol } ^ { - 1 }\)
B)239.9 kJ \(\mathrm { mol } ^ { - 1 }\)
C)998.3 kJ \(\mathrm { mol } ^ { - 1 }\)
D)2398 kJ \(\mathrm { mol } ^ { - 1 }\)
Q3) Define chemical energy.
Q4) Give the temperature and pressure for the standard state for a liquid.
Q5) Explain the difference between H and U.
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Chapter 7: The Quantum-Mechanical Model of the Atom
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Sample Questions
Q1) Identify the colour of a flame test for potassium.
A)violet
B)red
C)white
D)yellow
E)blue
Q2) Calculate the de Broglie wavelength of an egg with a mass of 60 g travelling at 32.18 m s<sup>-1</sup>.
A)5.0 × 10<sup>-34</sup> m
B)2.8 × 10<sup>-34</sup> m
C)4.9 × 10<sup>-34</sup> m
D)1.2 × 10<sup>-34</sup> m
E)3.4 × 10<sup>-34</sup> m
Q3) Define diffraction.
Q4) How many different values of m<sub>l</sub> are possible in the 2p sublevel? A)2
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Chapter 8: Periodic Properties of the Elements
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Sample Questions
Q1) How many valence electrons does an atom of Cu possess?
A)2
B)9
C)11
D)3
E)1
Q2) Identify a characteristic of halogens.
A)powerful reducing agents
B)forms water in reactions
C)powerful oxidizing agents
D)absorbs water in reactions
E)inert
Q3) How many valence electrons do the noble gases possess?
A)1
B)3
C)7
D)6
E)8
Q4) List the noble gas that has the highest ionization energy.
Q5) Define ionization energy.
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Chapter 9: Chemical Bonding I: Lewis Theory
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Sample Questions
Q1) Place the following in order of decreasing magnitude of lattice energy.
NaF \(\quad\)RbBr \(\quad\)KCl
A)RbBr > NaF > KCl
B)NaF > KCl > RbBr
C)KCl > NaF > RbBr
D)NaF > RbBr > KCl
E)RbBr > KCl > NaF
Q2) Give the number of lone pairs for BF<sub>3</sub>.
A)16
B)8
C)14
D)10
E)9
Q3) The phosphorus atom in PCl<sub>3</sub> would be expected to have a A)partial positive ( +)charge.
B)partial negative ( -)charge.
C)3+ charge.
D)3- charge.
Q4) Describe the difference between a pure covalent bond and a polar covalent bond.
Q5) Describe a covalent bond.
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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 have sp<sup>3</sup><sup>d</sup> hybridization on the central atom?
SiCl<sub>4</sub> \(\quad\)BrF<sub>5</sub> \(\quad\)AsF<sub>5</sub>\(\quad\) BrF<sub>3</sub>
A)2
B)0 C)4
D)1
E)3
Q2) In molecular orbital theory the acronym HOMO stands for A)homogenous orbital-molecular orientation.
B)highly ordered molecular orbital.
C)highest occupied molecular orbital.
D)highest orbital-molecular orientation.
E)It has no meaning.
Q3) Determine the hybridization about each interior atom in the following structure.Sketch the three-dimensional structure and label the interior atoms with their corresponding hybridization.
CH<sub>2</sub>CHCCCH<sub>3</sub>
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) Which of the following substances should have the highest melting point?
A)CO<sub>2</sub>
B)SrS
C)Xe
D)F<sub>2</sub>
E)MgO
Q2) Choose the molecule or compound that exhibits dispersion forces as its strongest intermolecular force.
A)Cl<sub>2</sub>
B)NH<sub>3</sub>
C)HF
D)NaCl
Q3) Place the following substances in order of increasing vapour pressure at a given temperature.
SF<sub>6</sub> \(\quad\)SiH<sub>4</sub> \(\quad\)SF<sub>4</sub>
A)SF<sub>6</sub> < SiH<sub>4 </sub>< SF<sub>4</sub>
B)SiH<sub>4</sub> < SF<sub>4</sub> < SF<sub>6</sub>
C)SF<sub>6 </sub>< SF<sub>4 </sub>< SiH<sub>4</sub>
D)SF<sub>4</sub> < SF<sub>6</sub> < SiH<sub>4</sub>
E)SiH<sub>4</sub> < SF<sub>6 </sub>< SF<sub>4</sub>
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Chapter 12: Solutions
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Sample Questions
Q1) How many grams of KBr are required to make 350.mL of a 0.115 mol L<sup>-1</sup> KBr solution?
A)0.338 g
B)3.04 g
C)4.79 g
D)40.3 g
Q2) Determine the solubility of N<sub>2</sub> in water exposed to air at 25 °C if the atmospheric pressure is 1.2 bar.Assume that the mole fraction of nitrogen is 0.78 in air and the Henry's law constant for nitrogen in water at this temperature is 6.1 × 10<sup>-4</sup> mol L<sup>-1 </sup>bar<sup>-1</sup>.
A)1.5 × 10<sup>-4</sup> mol L<sup>-1</sup>
B)6.5 × 10<sup>-4</sup> mol L<sup>-1</sup>
C)5.7 × 10<sup>-4</sup> mol L<sup>-1</sup>
D)1.8 × 10<sup>-4</sup> mol L<sup>-1</sup>
E)3.6 × 10<sup>-4</sup> mol L<sup>-1</sup>
Q3) A solution of LiCl in water has X<sub>LiCl</sub> = 0.0900.What is the molality?
A)4.46 mol kg<sup>-1</sup> LiCl
B)5.00 mol kg<sup>-1</sup> LiCl
C)5.49 mol kg<sup>-1</sup> LiCl
D)9.89 mol kg<sup>-1</sup> LiCl
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Chapter 13: Chemical Kinetics
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Sample Questions
Q1) Given the following rate law,how does the rate of reaction change if the concentration of X is doubled?
Rate = k [X][Y]<sup>2</sup>
A)The rate of reaction will increase by a factor of 2.
B)The rate of reaction will increase by a factor of 4.
C)The rate of reaction will increase by a factor of 5.
D)The rate of reaction will decrease by a factor of 2.
E)The rate of reaction will remain unchanged.
Q2) What is the unit of k in a zero-order reaction?
A)mol L<sup>-1</sup> s<sup>-1</sup>
B)mol L<sup>-1</sup> s
C)L mol<sup>-1 </sup>s<sup>-1</sup>
D)M<sup>2</sup> s<sup>-1</sup>
E)s<sup>-1</sup>
Q3) Identify a homogeneous catalyst.
A)SO<sub>2</sub> over vanadium (V)oxide
B)Pd in H<sub>2</sub> gas
C)Pt with methane
D)H<sub>2</sub>SO<sub>4</sub> with concentrated HCl
E)N<sub>2</sub> and H<sub>2</sub> catalyzed by Fe
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Chapter 14: Chemical Equilibrium
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Q1) Which of the following statements is TRUE?
A)If Q < K,it means the reverse reaction will proceed to form more reactants.
B)If Q > K,it means the forward reaction will proceed to form more products.
C)If Q = K,it means the reaction is at equilibrium.
D)If Q << K,it means the reaction is shifted to the left.
E)If Q >> K,it means the reaction is shifted to the right.
Q2) For the following reaction,what is n required in the conversion of K<sub>c</sub> to K<sub>p</sub>? \(\quad\)N<sub>2</sub>O<sub>4</sub>(g) 2NO<sub>2</sub>(g)
A)3
B)-1
C)-2
D)2
E)1
Q3) Identify the change that will always shift the equilibrium to the right.
A)remove reactant
B)increase product
C)remove product
D)increase pressure
E)increase volume
Q4) Why is a thermodynamic equilibrium constant unitless?
Page 16
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Chapter 15: Acids and Bases
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Sample Questions
Q1) The acid-dissociation constant of hydrocyanic acid (HCN)at 25.0 °C is 4.9 × \(10 ^ {10 }\) .What is the pH of an aqueous solution of 0.87 mol L<sup>-1</sup> sodium cyanide (NaCN)?
A)0.06
B)13.94
C)4.2 × 10<sup>-3</sup>
D)2.38
E)11.62
Q2) Describe a molecule that can be a Lewis acid.
Q3) Determine the pOH of a 0.116 mol L<sup>-1</sup> Ba(OH)<sub>2</sub> solution at 25 °C.
A)8.62
B)13.06
C)13.37
D)0.63
E)12.56
Q4) What is the difference between a strong and weak acid?
Q5) Describe the relationship between molecular structure and acid strength.
Q6) What is the autoionization of water?
Q7) What does the term amphoteric mean?
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Chapter 16: Aqueous Ionic Equilibrium
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Sample Questions
Q1) Calculate the percent ionization of nitrous acid in a solution that is 0.249 mol L<sup>-1</sup> in nitrous acid.The acid dissociation constant of nitrous acid is \(4.50 \times 10 ^ { - 4 }\)
A)5.53 %
B)0.0425 %
C)1.12 × \(10 ^ { - 4 }\) %
D)4.25 %
E)0.0450 %
Q2) A sample contains Ba<sub>3</sub>(PO<sub>4</sub>)<sub>2</sub>,<sub> </sub>CdS,AgCl,NH<sub>4</sub>Cl,and ZnS.Identify the precipitate after the addition of 6 mol L<sup>-1</sup> HCl; H<sub>2</sub>S and 0.2 mol L<sup>-1</sup> HCl; OH<sup></sup>to a pH of 8; and (NH<sub>4</sub>)<sub>2</sub>HPO<sub>4</sub> with NH<sub>3</sub>.
A)Ba<sub>3</sub>(PO<sub>4</sub>)<sub>2</sub> B)CdS
C)AgCl
D)NH<sub>4</sub>Cl E)ZnS
Q3) Predict the number of equivalence point(s)in the titration for a diprotic acid with a strong base.
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Chapter 17: Gibbs Energy and Thermodynamics
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Q1) Use the following thermodynamic values to calculate ?<sub>r</sub> <sub> </sub>\(G ^ { 0 }\) . \(\Delta_{r}\)\( H^{\circ} \) = +95 kJ \(\mathrm { mol } ^ { - 1 }\) ,\(\Delta_{r}\)\(S ^ { 0 }\) = -157 J \(\mathrm { mol } ^ { - 1 }\) ,T = 398 K
A)-48 kJ \(\mathrm { mol } ^ { - 1 }\)
B)-68 kJ \(\mathrm { mol } ^ { - 1 }\)
C)+39 kJ \(\mathrm { mol } ^ { - 1 }\)
D)+157 kJ \(\mathrm { mol } ^ { - 1 }\)
E)+142 kJ \(\mathrm { mol } ^ { - 1 }\)
Q2) Determine the equilibrium constant for the following reaction at 298 K: \(\quad\)Cl(g)+ O<sub>3</sub>(g)
ClO(g)+ O<sub>2</sub>(g)\(\quad\)\(\quad\) <sub>r</sub>G° = - 34.5 kJ mol<sup>-1</sup>
A)5.66 × 10<sup>5</sup>
B)0.986
C)8.96 × 10<sup>-7</sup>
D)4.98 × 10<sup>-4</sup>
E)1.12 × 10<sup>6</sup>
Q3) Why is heating your home with gas more efficient than heating it with electricity?
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Chapter 18: Electrochemistry
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Q1) How many electrons are transferred in the following reaction? (The reaction is unbalanced.)
\(\quad\)\(\quad\)Fe<sup>2+</sup>(aq)+ K(s) Fe(s)+ K<sup>+</sup>(aq)
A)1
B)2
C)3
D)4
E)6
Q2) Use the standard half-cell potentials listed below to calculate the standard cell potential for the following reaction occurring in an electrochemical cell at 2 5°C.(The equation is balanced.)
\(\quad\)\(\quad\)Pb(s)+ Br<sub>2</sub>(l) Pb<sup>2+</sup>(aq)+ 2 Br (aq) Pb<sup>2+</sup>(aq)+ 2 e<sup>-</sup> Pb(s)\(\quad\)\(\quad\) E° = -0.13 V Br<sub>2</sub>(l)+ 2 e<sup>-</sup> 2 Br<sup>-</sup> (aq)\(\quad\)\(\quad\)E° = +1.07 V
A)+1.20 V
B)+0.94 V
C)-0.94 V
D)-1.20 V
E)-0.60 V
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Chapter 19: Radioactivity and Nuclear Chemistry
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Q1) Neptunium-239 has a half-life of 2.35 days.How many days must elapse for a sample of <sup>239</sup>Np to decay to 0.100% of its original quantity?
A)0.0427 days
B)1.491 days
C)2.04 days
D)23.4 days
Q2) Determine how many neutrons are produced during the neutron-induced fission of 239
94 Pu to form 89
36 Kr and 149
58 Ce.
A)2
B)0
C)3
D)1
E)4
Q3) Describe what is meant by the "valley of stability."
Q4) List two problems associated with nuclear power.
Q5) What is the mass defect?
Q6) How does a dosimeter measure exposure to radioactivity?
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Chapter 20: Organic Chemistry I: Structures
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Q1) Which of the following organic families contains nonpolar compounds?
A)carboxylic acids
B)alcohols
C)amines
D)esters
E)ethers
Q2) What is the index of hydrogen deficiency in C<sub>9</sub>H<sub>16</sub>Br<sub>2</sub>?
A)5
B)4
C)3
D)2
E)1
Q3) What does the term "racemic mixture" mean?
Q4) Which of the following condensed general formulas represents tertiary alcohols?
A)RCH<sub>2</sub>OR
B)RCH<sub>2</sub>OH
C)R<sub>3</sub>COR
D)R<sub>3</sub>COH
E)R<sub>3</sub>CCH<sub>2</sub>OH
22
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Chapter 21: Organic Chemistry II: Reactions
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Q1) Complete and balance the following addition reaction: CH<sub>3</sub>C CCH<sub>3</sub> + 2Cl<sub>2</sub> ?
A)CH<sub>3</sub>C CCH<sub>3</sub> + 2Cl<sub>2</sub>
CH<sub>3</sub>CCl<sub>2</sub>CCl<sub>2</sub>CH<sub>3</sub>
B)CH<sub>3</sub>C CCH<sub>3</sub> + 3Cl<sub>2</sub>
2CH<sub>3</sub>CCl<sub>3</sub>
C)CH<sub>3</sub>C CCH<sub>3</sub> + 2Cl<sub>2</sub> 4CH<sub>3</sub>Cl
D)CH<sub>3</sub>C CCH<sub>3</sub> + 4Cl<sub>2</sub> 4CH<sub>2</sub>Cl<sub>2</sub>
E)CH<sub>3</sub>C CCH<sub>3</sub> + Cl<sub>2</sub>
CH<sub>3</sub>CHClCHClCH<sub>3</sub>
Q2) Choose the weak acid from the compounds below.
A)CH<sub>3</sub>CH<sub>2</sub>NH<sub>2</sub>
B)CH<sub>3</sub>CH<sub>2</sub>CO<sub>2</sub>H
C)CH<sub>3</sub>CH<sub>2</sub>OCH<sub>3</sub>
D)CH<sub>3</sub>CH<sub>2</sub>F
E)CH<sub>3</sub>COOC<sub>2</sub>H<sub>5</sub>
Q3) Draw a clear mechanism for E2 elimination reactions.Make sure to indicate each step of the mechanism (if there is more than one)as well as to use arrows to show the electron flow.
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Chapter 22: Biochemistry
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Q1) Which one of the following amino acids contains a polar side chain?
A)alanine
B)glycine
C)threonine
D)valine
Q2) Which of the following shows all of the tripeptides that can be formed from one molecule each of lysine (Lys),threonine (Thr),and histidine (His)?
A)LysThrHis
B)LysThrHis,LysHisThr,HisThrLys,HisLysThr,ThrLysHis,ThrHisLys
C)LysThrHis,LysHisThr,HisLysThr,ThrHisLys
D)LysThrHis,LysHisThr,HisLysThr
E)LysThrHis and LysHisThr
Q3) Identify the polysaccharide.
A)glucose
B)cholesterol
C)fructose
D)cellulose
E)sucrose
Q4) Why does the melting point of fatty acids decrease with the number of double bonds present?
Page 24
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Chapter 23: Chemistry of the Nonmetals
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Q1) What is the oxidation number of oxygen in Na<sub>2</sub>O? A)-4
Q2) Determine the number of vertices and faces in the closo-Borane B<sub>10</sub>H<sub>10</sub><sup>2-</sup>.
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
Q3) Refer to the above table.Use charge balance to determine the value of x in the pyrosilicate Zn<sub>x</sub>Si<sub>2</sub>O<sub>7</sub>.
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Chapter 24: Metals and Metallurgy
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Q1) What is the difference between ferromagnetism and paramagnetism?
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) Identify the normal crystal structure at room temperature for Ni.
A)face-centred cubic
B)body-centred cubic
C)tetrahedral
D)hexagonal
E)hexagonal closest packed
Q4) What is the highest oxidation number of chromium?
A)+3
B)+4
C)+5
D)+7
E)+6
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Chapter 25: Transition Metals and Coordination Compounds
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Q1) How many d electrons are there in Cr<sub>2</sub>O<sub>7</sub><sup>2-</sup>?
A)0
B)1
C)2
D)3
Q2) Choose the electron configuration for Cr<sup>3+</sup>.
A)[Ne]3s<sup>2</sup>3p<sup>6</sup>
B)[Ar]4s<sup>1</sup>3d<sup>2</sup>
C)[Ar]4s<sup>2</sup>3d<sup>1</sup>
D)[Ar]4s<sup>2</sup>3d<sup>7</sup>
E)[Ar]3d<sup>3</sup>
Q3) Determine the chemical formula for the compound tetracarbonylplatinum(IV)chloride.
A)[Pt(CO)<sub>4</sub>]Cl<sub>2</sub>
B)[Pt(CO)<sub>4</sub>]Cl<sub>4</sub>
C)[Pt(CO)<sub>4</sub>Cl<sub>4</sub>]<sup>2-</sup>
D)[PtCl<sub>4</sub>](CO)<sub>4</sub>
E)[Pt(CO)<sub>4</sub>Cl<sub>4</sub>]<sup>4-</sup>
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?