Answers to Even Chapter 6 Review Questions PDF

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First year chemistry 110 Practise questions and skill drills...

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CHAPTER 6 Chemistry of Water, Chemistry in Water REVIEW QUESTIONS Section 6.1: Arsenic Ain’t Arsenic 6.50 (a) H2CO3(aq) molecules, HCO3−(aq) ions, CO32−(aq) ions (b) For example, Fe2+(aq) ions, Fe3+(aq) ions, [Fe(CN)6]3−(aq) ions. Note that the first two ions are better described as [Fe(OH2)6]2+(aq) and [Fe(OH2)6]3+(aq)—especially the latter for which the “coordination bonds” between Fe3+ and O (of H2O) are stronger. Also, depending upon the pH, there may be significant amounts of [Fe(OH)(OH2)5]2+(aq) ions coexisting with [Fe(OH2)6]3+(aq) ions. [Fe(OH2)6]3+(aq) ion is a weak acid.

6.52 Arsenous acid (H3AsO3 or As(OH)3) is a trivalent, pyramidal, polyprotic acid. This can be categorized as an inorganic arsenic compound. When dissolved in water, its ions dissociate for a net reaction of:

H3AsO3 (aq) + 3 H2O(

H3O] (aq) + [AsO 3] 3(aq)



Methanearsonic acid is a pentavalent, tetrahedral, diprotic acid. This can be categorized as an organic arsenic compound. When dissolved in water, its ions dissociate for a net reaction of:

MeAsO(OH) 2 (aq) + 2 H 2O(



H 3O] (aq) + [MeAsO 3] 2(aq)

Arsenite ions ([AsO3]3‒) are known to interact with sulfur containing groups on certain enzymes. And while not all organic arsenic aqueous ions are non-poisonous, they can be less harmful than inorganic aqueous ions and are excreted instead of being metabolized. Therefore, an aqueous solution of methanearsonic acid would be less toxic to humans.

6.54 H2 O( (a) Ca(NO3 )2 (s)  Ca2  (aq) 2NO3  (aq) Ca2+(aq) is present in lower concentrations H O(

2

2

2 (b) MgCO3 (s)   Mg  (aq)  CO3  (aq) Both have equal concentrations

Chapter 6: Chemistry of Water, Chemistry in Water

6.56 H2O is the chemical formula. H2O(g) denotes water in the gaseous state. H2O(ℓ) denotes water in the liquid state. Section 6.2: The Remarkable Properties of Water 6.58 The paper clip can float on water even though it is made of a metal denser than water because of surface tension. In order for the paper clip to sink to the bottom of the water, it must first distort the surface water and temporarily increase the surface area. Because surface tension is positive, this represents an energy barrier that can keep the clip from sinking. If you do not disturb the paper clip too much—i.e., give it too much energy—while setting it down, you can get it to float. It is less likely that a paper clip can float on the surface of octane as there is less positive surface tension, the surface is easily distorted and the paperclip will sink.

Section 6.3: Intermolecular Forces 6.60 (a) Si→O is a polar bond, whereas P‒P is non-polar (b) C→O and C→S are both polar bonds (C),  (S) and (O) = 2.5, 2.6 and 3.5 C→O has the greater polarity 6.62 (a) All of the bonds in urea, N‒H, N‒C, and C=O, are polar. (b) The C=O bond is the most polar bond in urea; Δχ = |χ(O) − χ(C)| = |3.5 − 2.5| = 1. O is the negative end of the bond, i.e., the dipole looks like C→O.

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Chapter 6: Chemistry of Water, Chemistry in Water 6.64

(a) Δχ = 1.3 in bent H2O, 0.8 in ClF, and 0.8 in pyramidal NH3. H2O > ClF > NH3 is the predicted order of polarities. Note that while the two H‒O bond dipoles partially cancel in H2O, the net dipole is still considerable because there are two bonds rather than one as in ClF. Note that Δχ is the same for both Cl‒F and N‒H bonds. However, the three N‒H dipoles partially cancel in NH3, making it less polar than ClF. (b) Linear CO2 and tetrahedral CCl4 are not polar. Though they have polar bonds, the bond dipoles cancel due to symmetry. (c) F in ClF is more negatively charged.

6.66 (a) CO is polar. (b) Trigonal planar BCl3 is non-polar due to symmetry. (c) Tetrahedral CF4 is non-polar due to symmetry. (d) Trigonal pyramidal PCl3 is polar. (e) Tetrahedral GeH4 is non-polar due to symmetry. Δχ = 1.5 in CF4, 1.2 in BCl3, and 1.0 in pyramidal PCl3. The most polar bonds are in CF4, though the molecule has no net dipole. 6.68 (b), (e), (d), (g) and (h) i.e., hydrogen bonds occur in liquid NH3, CH3COOH, HF, HOCH2CH2OH, and CH3NH2. There are no hydrogen bonds in liquid C2H5OC2H5, CH4, or Br2. 6.70

An acetic acid dimer 6.72 (a) When ice melts, H-bonds (special dipole‒dipole interactions) and dispersion forces are overcome. (b) When solid I2 sublimes, dispersion forces are overcome.

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Chapter 6: Chemistry of Water, Chemistry in Water (c) When NH3 passes from liquid to vapour phase, H-bonds (special dipole‒dipole interactions) and dispersion forces are overcome. 6.74 (a) Dispersion forces operate in liquid O2. (b) Hydrogen bonding and dispersion forces operate in methanol. (c) Dipole‒dipole and dispersion forces operate in liquid sulfur dioxide, SO2, whose molecules are V-shaped. (d) Dispersion forces operate in linear carbon dioxide, CO2(s). 6.76 Polarity refers to a positive and negative charge separation, i.e., a dipole, in a molecule. Polarizability refers to the response of a molecule to an electric field. When a molecule is exposed to an electric dipole, it will develop an electric dipole in the opposite direction. (− +) ← molecule with dipole next to (δ+ δ−) ← polarizable molecule with induced dipole Compounds with a high polarity have higher boiling points. 6.78 Hydrogen bonding in liquid water is depicted below (solid line: covalent bond. dashed line: hydrogen bonding interaction, dots: lone pairs of electrons).

O H

O H

H

H O H

H

O H

O H

H

H

6.80 Sodium chloride is an ionic substance. χCl >> χNa (3.2 vs. 0.9). ∆χ = 3.2 ‒ 0.9 = 2.3. The bonding pair of electrons shifts toward the more electronegative atom, Cl, creating a dipole. Therefore, NaCl is a polar substance.

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Chapter 6: Chemistry of Water, Chemistry in Water Section 6.4: Explaining the Properties of Water 6.82 The boiling points of H2O and H2S are 100°C and ‒60°C, respectively (water has the higher boiling point). This can be due to 2 factors: (i) From an electronegativity standpoint: χ(O) > χ(S). For O‒H bonds, ∆χ = 3.5 ‒ 2.2 = 1.3. For S‒H bonds, ∆χ = 2.6 ‒ 2.2 = 0.4. Indeed, O‒H is the stronger bond. (ii) H2S molecules interact with themselves through London forces (dispersion forces) and dipole‒dipole interactions. H2O has the added benefit of hydrogen-bonding interactions in conjunction with both of the aforementioned forces, due to the extremely electronegative oxygen atom pulling the adjacent molecules toward it. 6.84 The trend in the group 16 H2X compounds is for the boiling point to increase as X descends the group. H2O does not fit this trend. The pattern is due to the increasing dispersion forces as the X atom gets bigger. H2O does not fit the trend because hydrogen bonds (especially strong dipole‒dipole interactions) form between H and O atoms on neighbouring molecules. These forces are so strong that H2O even has a higher boiling point than H2Te. S atoms are not sufficiently electronegative, and are too big for hydrogen bonding to operate in H2S. 6.86 Ethanol, CH3CH2OH(ℓ), has a higher boiling point than dimethyl ether, CH3OCH3(ℓ), a compound with the same compositional formula C2H6O, because ethanol molecules can form hydrogen bonds, whereas dimethyl ether cannot. Both molecules have dipole‒dipole forces, but they are the especially strong hydrogen bonds in the case of ethanol. 6.88 (a) In methanol, hydrogen bonding and dispersion forces affect the enthalpy change of vaporization, vapH (b) In octane, only dispersion forces affect the enthalpy change of vaporization, vapH (c) In acetone, dipole-dipole and dispersion forces affect the enthalpy change of vaporization, vapH 6.90 (a) In cis-1,2-dichloroethylene dipole‒dipole and dispersion forces affect the equilibrium vapour pressure. In this case, the C‒Cl bond dipoles reinforce each other (partially) because they are on the same side of the molecule. The molecule has a net dipole moment. (b) In trans-1,2-dichloroethylene dispersion forces affect the equilibrium vapour pressure. In this case, the C‒Cl bond dipoles cancel out because they point in opposite directions. trans-1,2-dichloroethylene should have the higher vapour pressure at a given temperature.

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Chapter 6: Chemistry of Water, Chemistry in Water 6.92 Liquid ethylene glycol, HOCH2CH2OH, should have greater viscosity than ethanol, CH3CH2OH, because it has two O-H groups, versus one for ethanol, and can form more hydrogen bonds. Maleic acid, on the left, makes a strong intramolecular hydrogen bond; this reduces opportunities for intermolecular hydrogen bonds, as an O and H are already hydrogen bonding. Strong intermolecular pairs of hydrogen bonds are formed between adjacent fumaric acid molecules. Section 6.5: Water as a Solvent 6.94 Sodium nitrate (NaNO3) is soluble, while CaCO3 is not. The degree of solubility for ionic substances is the result of competitive anion‒cation attractions in the lattice and ion‒dipole attractions to water molecules surrounding the ions. Consult Figure 6.28 as a guideline to predict solubility of ionic compounds. From this table, we see that the oxidation state for the salts of ions that are typically soluble exist in either +1 or ‒1, with some exceptions. 6.96 (a) (b) (c) (d)

The O end (the negative end) of water points to Ca2+(aq). The H end (the positive end) of water points to Br−(aq). The H end (the positive end) of water points to Cr2O72−(aq). The O end (the negative end) of water points to NH4+(aq).

6.98 When silver nitrate, AgNO3(s), and potassium chloride, KCl(s), are dissolved in some water, the possible reactant species are H2O(ℓ), Ag+(aq) ions, NO3−(aq) ions, K+(aq) ions, and Cl−(aq) ions. In fact, Ag+(aq) ions and Cl−(aq) ions are the reactants in a precipitation reaction. The other two solvated ions are spectators: they do not participate in the reaction.

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Chapter 6: Chemistry of Water, Chemistry in Water 6.100 Cooking oil is not miscible with water because its molecules are non-polar (or very weakly polar). They do not interact with water strongly enough for water to solvate them; water molecules prefer to interact with other water molecules. Cooking oil is soluble in hexane; a nonpolar solvent.

Section 6.6: Self-Ionization of Water 6.102 The concentrations of hydronium ions and hydroxide ions are equal in pure water because of the stoichiometry of the self-ionization reaction, 2 H2O(ℓ) → H3O+(aq) + OH−(aq) i.e., we get one hydronium ion for each hydroxide ion formed.

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Chapter 6: Chemistry of Water, Chemistry in Water Section 6.7: Categories of Chemical Reaction in Water 6.104 Since barium sulfate, BaSO4(s), precipitates from water, whereas iron (II) sulfate does not, we conclude that BaSO4(s) is insoluble in water (it has a very low solubility) whereas FeSO4(s) is soluble. 6.106 2 Cr2O 72  (aq) + 28 H + (aq) + 12 e   4 Cr3+ (aq) + 14 H 2 O( 3C2 H 5OH(aq) + 3 H 2 O(   3CH3 COOH(aq) + 12 H+ (aq) + 12 e Here, ethanol, C2H5OH(aq), is oxidized to acetic acid, CH3COOH(aq). Aquated dichromate ions, Cr2O72−(aq), are the oxidizing agent. Cr2O72−(aq) ions are reduced to Cr3+(aq) ions. Ethanol is the reducing agent.

6.108 When nitric acid dissolves in water, nitrate and hydronium ions are produced. Because nitric acid is a strong acid, there is a negligible concentration of undissociated aquated HNO3. When barium hydroxide dissolves in water, Ba2+(aq) ions and OH−(aq) ions are produced. 6.110 (a) An iron(II) ion is bound to six ammonia molecules: [Fe(NH3)6]2+

The net charge on this complex is +2.

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Chapter 6: Chemistry of Water, Chemistry in Water (b) A zinc ion is bound to four cyanide ions: [Zn(CN)4]2-

The net charge is −4 + 2 = −2. (c) A manganese(II) ion is bound to six fluoride ions: [MnF6]4-

The net charge is −6 + 2 = −4. (d) An iron(III) ion is bound to six cyanide ions: [Fe(CN)6]3-

The net charge is −6 + 3 = −3.

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Chapter 6: Chemistry of Water, Chemistry in Water (e) A cobalt(II) ion is bound to four chloride ions: [CoCl4]2-

The net charge is −4 + 2 = −2. (f) A nickel(II) ion is bound to four ammonia molecules and two water molecules: [Ni(NH3)4(OH2)2]2+

This is the trans isomer. The net charge is +2

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Chapter 6: Chemistry of Water, Chemistry in Water Section 6.8: Solution Concentration 6.112 H O(

2

2 BaCl 2 (s)   Ba  (aq)  2Cl (aq)

The label on the flask should be 0.15 mol L‒1 BaCl2 solution.

6.114 (a) (b) (c) (d) (e)

(A) 0.20 mol sucrose > 0.05 mol sucrose (B) 0.50 mol sucrose > 0.20 mol sucrose (A) and (B) both have 0.20 mol sucrose (B) 0.20 mol sucrose > 0.020 mol sucrose (A) in this case, (B) is just 1/20th of (A), i.e., it has 1/20th the amount of sucrose. There is no 6.116, 6.118 or 6.120 in the textbook … visualize a shoulder shrug here …

6.122 (a) 0.12 mol L−1 Ba2+(aq) ions and 0.24 mol L−1 Cl−(aq) ions (b) 0.0125 mol L−1 Cu2+(aq) ions and SO42−(aq) ions (c) 1.00 mol L-1 K+(aq) ions and 0.500 mol L−1 Cr2O72−(aq) ions 6.124 Benzoic acid is a weak acid. From Table 6.13 (Section 6.7), we know that acetic acid is also a weak acid. The text states that only 1 to 5% of acetic acid molecules are ionized in dilute solutions. If we apply this to a 1.00 mol L‒1 solution, the approximate concentrations for the ions at maximum ionization would be [CH3COO‒] = [H+] = 0.05 mol L‒1, [CH3COOH] = 0.95 mol L‒1. Based on this argument, it appears that acetic acid is a stronger acid than benzoic. However, consulting the pKa at the back of the text, it is clear that benzoic acid (4.2) is a stronger acid than acetic acid (4.74). SUMMARY AND CONCEPTUAL QUESTIONS 6.126 The pressure of water vapour in the space above the liquid in flask A is the same as in flask B. Both flasks are in a state of equilibrium vapour pressure, where molecules entering the headspace (entering the vapour phase) are occurring simultaneously with those returning to the liquid phase. At no point is there a further increase in vapour pressure. Water in both flasks will exhibit a vapour pressure of 3.17 kPa regardless of volume in the sealed flask, unless the temperature is increased.

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Chapter 6: Chemistry of Water, Chemistry in Water

6.128 (a) Reduction: “what happens to a reactant species that takes electrons from another” vs. “something made smaller” (b) Complex: “product formed through ligand binding” vs. “consisting of many parts” (c) Strong: “dissociates completely” vs. “exerting great force” (d) Spectator: “species that does not participate in a reaction” vs. “person who watches event, sports, show, etc. but does not participate” 6.130 (a) A sodium atom is highly electropositive (χ = 0.9) and readily gives up its valence electron so that it can become stable with a full valence shell of 8 electrons, forming a sodium cation (Na+). (b) Like sodium, magnesium is electropositive (χ = 1.3) and would like to give up its two valence electrons to have a full shell of 8 electrons to form a magnesium dication (Mg2+). Oxygen is electronegative (χ = 3.5) and a perfect receptor for these two electrons to form an O2‒ anion. (c) Electronegativity increases across the period and up. Group 17 is home to the most electronegative atoms, including the most electronegative atom fluorine (χ = 4.0). However, chlorine is less electronegative than oxygen since it is in the third period and its electrons are more shielded. (d) Ionic compounds are formed when electropositive atoms (M) “transfer” or “donate” their n valence electrons to X atoms in order to become stable with a full valence shell.

6.132

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The figure on the left is a frame from the animation in e6.4 molecular-level modelling activity depicting the inside of a bubble in boiling water. (a) Kinetic energy of the water molecules within and on the surface of the bubble prevents its collapse. These “hot” molecules smash into their neighbours pushing open the bubble. (b) Bubbles form when the vapour pressure within the heated liquid slightly exceeds atmospheric pressure— enough to push open the bubble.

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Chapter 6: Chemistry of Water, Chemistry in Water

6.134

The figure on the left is a frame from the animation in e6.21 molecular-level modelling activity depicting a hydronium ion transferring a proton to another water molecule. (a) Hydronium ions appear to diffuse very quickly through an aqueous solution—faster than they should be able to via diffusion—because the transfer of hydronium across a cluster of water molecules is achieved through successive H+ ion transfers between adjacent water molecules. (b) Each proton transfer involves one O atom of a hydronium ion taking the pair of electrons from an H atom bonded to it, and another O atom (on a water molecule on the other side of the H) donating a lone pair of electrons to the released proton—i.e., accepting the proton.

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Chapter 6: Chemistry of Water, Chemistry in Water

6.136

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The figure on the left is a frame from the animation in e6.22 molecular-level modelling activity depicting the formation of a AgCl precipitate. (a) Here, the electrostatic attraction of the Ag+ and Cl− ions overcomes the dipole-ion forces between solvating water molecules and the ions to cause growth of the ionic lattice. (b) The lattice is an array of Ag+ and Cl− ions. It does not consist of distinct AgCl molecules.

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Chapter 6: Chemistry of Water, Chemistry in Water

6.142 (a) Ca(OH)2 (s) + 2 H3O+ (aq)   Ca 2+ (aq) + 4 H 2O( is an acid-base reaction, (iii). + (b) CH3 COOH(aq) + Ag+ (aq) + H2O(l)   AgCH3COO(s) + H 3O (aq) is a precipitation reaction, (i). (c) Fe(OH) 3(s) + 3H 2C 2O 4(aq)   [Fe(C 2O 4) 3]3 (aq) + 3H 2O( is an acid-base, (iii), and a complexation, (iv), reaction. The oxalic acid neutralizes the hydroxyls initially coordinated to Fe2+ (an acid-base reaction), then forms a bond to the iron ion itself—it is a bidentate ligand. (d) 5Fe2+ (aq) + MnO4 (aq) + 8 H+ (aq)   5Fe3+ (aq) + Mn 2+ (aq) + 4H 2O( is an oxidation-reduction reaction, (ii). 6.143 The bond between a Lewis acid and a Lewis base is a covalent bond formed with both of the electrons coming from the Lewis base.

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