Private: Chapter Eighteen
Key Terms, Key Equations, Summaries, and Exercises (Chapter 18)
Key Terms
- acid anhydride
- compound that reacts with water to form an acid or acidic solution
- alkaline earth metal
- any of the metals (beryllium, magnesium, calcium, strontium, barium, and radium) occupying group 2 of the periodic table; they are reactive, divalent metals that form basic oxides
- allotropes
- two or more forms of the same element, in the same physical state, with different chemical structures
- amorphous
- solid material such as a glass that does not have a regular repeating component to its three-dimensional structure; a solid but not a crystal
- base anhydride
- metal oxide that behaves as a base towards acids
- bicarbonate anion
- salt of the hydrogen carbonate ion, HCO3−
- bismuth
- heaviest member of group 15; a less reactive metal than other representative metals
- borate
- compound containing boron-oxygen bonds, typically with clusters or chains as a part of the chemical structure
- carbonate
- salt of the anion CO32−; often formed by the reaction of carbon dioxide with bases
- chemical reduction
- method of preparing a representative metal using a reducing agent
- chlor-alkali process
- electrolysis process for the synthesis of chlorine and sodium hydroxide
- disproportionation reaction
- chemical reaction where a single reactant is simultaneously reduced and oxidized; it is both the reducing agent and the oxidizing agent
- Downs cell
- electrochemical cell used for the commercial preparation of metallic sodium (and chlorine) from molten sodium chloride
- Frasch process
- important in the mining of free sulfur from enormous underground deposits
- Haber process
- main industrial process used to produce ammonia from nitrogen and hydrogen; involves the use of an iron catalyst and elevated temperatures and pressures
- halide
- compound containing an anion of a group 17 element in the 1− oxidation state (fluoride, F−; chloride, Cl−; bromide, Br−; and iodide, I−)
- Hall–Héroult cell
- electrolysis apparatus used to isolate pure aluminum metal from a solution of alumina in molten cryolite
- hydrogen carbonate
- salt of carbonic acid, H2CO3 (containing the anion HCO3−) in which one hydrogen atom has been replaced; an acid carbonate; also known as bicarbonate ion
- hydrogen halide
- binary compound formed between hydrogen and the halogens: HF, HCl, HBr, and HI
- hydrogen sulfate
- HSO4− ion
- hydrogen sulfite
- HSO3− ion
- hydrogenation
- addition of hydrogen (H2) to reduce a compound
- hydroxide
- compound of a metal with the hydroxide ion OH− or the group −OH
- interhalogen
- compound formed from two or more different halogens
- metal (representative)
- atoms of the metallic elements of groups 1, 2, 12, 13, 14, 15, and 16, which form ionic compounds by losing electrons from their outer s or p orbitals
- metalloid
- element that has properties that are between those of metals and nonmetals; these elements are typically semiconductors
- nitrate
- NO3− ion; salt of nitric acid
- nitrogen fixation
- formation of nitrogen compounds from molecular nitrogen
- Ostwald process
- industrial process used to convert ammonia into nitric acid
- oxide
- binary compound of oxygen with another element or group, typically containing O2− ions or the group –O– or =O
- ozone
- allotrope of oxygen; O3
- passivation
- metals with a protective nonreactive film of oxide or other compound that creates a barrier for chemical reactions; physical or chemical removal of the passivating film allows the metals to demonstrate their expected chemical reactivity
- peroxide
- molecule containing two oxygen atoms bonded together or as the anion, O22−
- photosynthesis
- process whereby light energy promotes the reaction of water and carbon dioxide to form carbohydrates and oxygen; this allows photosynthetic organisms to store energy
- Pidgeon process
- chemical reduction process used to produce magnesium through the thermal reaction of magnesium oxide with silicon
- polymorph
- variation in crystalline structure that results in different physical properties for the resulting compound
- representative element
- element where the s and p orbitals are filling
- representative metal
- metal among the representative elements
- silicate
- compound containing silicon-oxygen bonds, with silicate tetrahedra connected in rings, sheets, or three-dimensional networks, depending on the other elements involved in the formation of the compounds
- sulfate
- SO42− ion
- sulfite
- SO32− ion
- superoxide
- oxide containing the anion O2−
Summaries
18.1 Periodicity
This section focuses on the periodicity of the representative elements. These are the elements where the electrons are entering the s and p orbitals. The representative elements occur in groups 1, 2, and 12–18. These elements are representative metals, metalloids, and nonmetals. The alkali metals (group 1) are very reactive, readily form ions with a charge of 1+ to form ionic compounds that are usually soluble in water, and react vigorously with water to form hydrogen gas and a basic solution of the metal hydroxide. The outermost electrons of the alkaline earth metals (group 2) are more difficult to remove than the outer electron of the alkali metals, leading to the group 2 metals being less reactive than those in group 1. These elements easily form compounds in which the metals exhibit an oxidation state of 2+. Zinc, cadmium, and mercury (group 12) commonly exhibit the group oxidation state of 2+ (although mercury also exhibits an oxidation state of 1+ in compounds that contain Hg22+). Aluminum, gallium, indium, and thallium (group 13) are easier to oxidize than is hydrogen. Aluminum, gallium, and indium occur with an oxidation state 3+ (however, thallium also commonly occurs as the Tl+ ion). Tin and lead form stable divalent cations and covalent compounds in which the metals exhibit the 4+-oxidation state.
18.2 Occurrence and Preparation of the Representative Metals
Because of their chemical reactivity, it is necessary to produce the representative metals in their pure forms by reduction from naturally occurring compounds. Electrolysis is important in the production of sodium, potassium, and aluminum. Chemical reduction is the primary method for the isolation of magnesium, zinc, and tin. Similar procedures are important for the other representative metals.
18.3 Structure and General Properties of the Metalloids
The elements boron, silicon, germanium, arsenic, antimony, and tellurium separate the metals from the nonmetals in the periodic table. These elements, called metalloids or sometimes semimetals, exhibit properties characteristic of both metals and nonmetals. The structures of these elements are similar in many ways to those of nonmetals, but the elements are electrical semiconductors.
18.4 Structure and General Properties of the Nonmetals
Nonmetals have structures that are very different from those of the metals, primarily because they have greater electronegativity and electrons that are more tightly bound to individual atoms. Most nonmetal oxides are acid anhydrides, meaning that they react with water to form acidic solutions. Molecular structures are common for most of the nonmetals, and several have multiple allotropes with varying physical properties.
18.5 Occurrence, Preparation, and Compounds of Hydrogen
Hydrogen is the most abundant element in the universe and its chemistry is truly unique. Although it has some chemical reactivity that is similar to that of the alkali metals, hydrogen has many of the same chemical properties of a nonmetal with a relatively low electronegativity. It forms ionic hydrides with active metals, covalent compounds in which it has an oxidation state of 1− with less electronegative elements, and covalent compounds in which it has an oxidation state of 1+ with more electronegative nonmetals. It reacts explosively with oxygen, fluorine, and chlorine, less readily with bromine, and much less readily with iodine, sulfur, and nitrogen. Hydrogen reduces the oxides of metals with lower reduction potentials than chromium to form the metal and water. The hydrogen halides are all acidic when dissolved in water.
18.6 Occurrence, Preparation, and Properties of Carbonates
The usual method for the preparation of the carbonates of the alkali and alkaline earth metals is by reaction of an oxide or hydroxide with carbon dioxide. Other carbonates form by precipitation. Metal carbonates or hydrogen carbonates such as limestone (CaCO3), the antacid Tums (CaCO3), and baking soda (NaHCO3) are common examples. Carbonates and hydrogen carbonates decompose in the presence of acids and most decompose on heating.
18.7 Occurrence, Preparation, and Properties of Nitrogen
Nitrogen exhibits oxidation states ranging from 3− to 5+. Because of the stability of the N≡N triple bond, it requires a great deal of energy to make compounds from molecular nitrogen. Active metals such as the alkali metals and alkaline earth metals can reduce nitrogen to form metal nitrides. Nitrogen oxides and nitrogen hydrides are also important substances.
18.8 Occurrence, Preparation, and Properties of Phosphorus
Phosphorus (group 15) commonly exhibits oxidation states of 3− with active metals and of 3+ and 5+ with more electronegative nonmetals. The halogens and oxygen will oxidize phosphorus. The oxides are phosphorus(V) oxide, P4O10, and phosphorus(III) oxide, P4O6. The two common methods for preparing orthophosphoric acid, H3PO4, are either the reaction of a phosphate with sulfuric acid or the reaction of water with phosphorus(V) oxide. Orthophosphoric acid is a triprotic acid that forms three types of salts.
18.9 Occurrence, Preparation, and Compounds of Oxygen
Oxygen is one of the most reactive elements. This reactivity, coupled with its abundance, makes the chemistry of oxygen very rich and well understood.
Compounds of the representative metals with oxygen exist in three categories (1) oxides, (2) peroxides and superoxides, and (3) hydroxides. Heating the corresponding hydroxides, nitrates, or carbonates is the most common method for producing oxides. Heating the metal or metal oxide in oxygen may lead to the formation of peroxides and superoxides. The soluble oxides dissolve in water to form solutions of hydroxides. Most metals oxides are base anhydrides and react with acids. The hydroxides of the representative metals react with acids in acid-base reactions to form salts and water. The hydroxides have many commercial uses.
All nonmetals except fluorine form multiple oxides. Nearly all of the nonmetal oxides are acid anhydrides. The acidity of oxyacids requires that the hydrogen atoms bond to the oxygen atoms in the molecule rather than to the other nonmetal atom. Generally, the strength of the oxyacid increases with the number of oxygen atoms bonded to the nonmetal atom and not to a hydrogen.
18.10 Occurrence, Preparation, and Properties of Sulfur
Sulfur (group 16) reacts with almost all metals and readily forms the sulfide ion, S2−, in which it has as oxidation state of 2−. Sulfur reacts with most nonmetals.
18.11 Occurrence, Preparation, and Properties of Halogens
The halogens form halides with less electronegative elements. Halides of the metals vary from ionic to covalent; halides of nonmetals are covalent. Interhalogens form by the combination of two or more different halogens.
All of the representative metals react directly with elemental halogens or with solutions of the hydrohalic acids (HF, HCl, HBr, and HI) to produce representative metal halides. Other laboratory preparations involve the addition of aqueous hydrohalic acids to compounds that contain such basic anions, such as hydroxides, oxides, or carbonates.
18.12Occurrence, Preparation, and Properties of the Noble Gases
The most significant property of the noble gases (group 18) is their inactivity. They occur in low concentrations in the atmosphere. They find uses as inert atmospheres, neon signs, and as coolants. The three heaviest noble gases react with fluorine to form fluorides. The xenon fluorides are the best characterized as the starting materials for a few other noble gas compounds.
Exercises
18.1 Periodicity
How do alkali metals differ from alkaline earth metals in atomic structure and general properties?
Why does the reactivity of the alkali metals decrease from cesium to lithium?
Predict the formulas for the nine compounds that may form when each species in column 1 of Exercise 18.3 reacts with each species in column 2.
| 1 | 2 |
|---|---|
| Na | I |
| Sr | Se |
| Al | O |
Predict the best choice in each of the following. You may wish to review the chapter on electronic structure for relevant examples.
(a) the most metallic of the elements Al, Be, and Ba
(b) the most covalent of the compounds NaCl, CaCl2, and BeCl2
(c) the lowest first ionization energy among the elements Rb, K, and Li
(d) the smallest among Al, Al+, and Al3+
(e) the largest among Cs+, Ba2+, and Xe
Sodium chloride and strontium chloride are both white solids. How could you distinguish one from the other?
The reaction of quicklime, CaO, with water produces slaked lime, Ca(OH)2, which is widely used in the construction industry to make mortar and plaster. The reaction of quicklime and water is highly exothermic:
CaO(𝑠)+H2O(𝑙)⟶Ca(OH)2(𝑠)Δ𝐻=−350 kJmol−1
(a) What is the enthalpy of reaction per gram of quicklime that reacts?
(b) How much heat, in kilojoules, is associated with the production of 1 ton of slaked lime?
Write a balanced equation for the reaction of elemental strontium with each of the following:
(a) oxygen
(b) hydrogen bromide
(c) hydrogen
(d) phosphorus
(e) water
How many moles of ionic species are present in 1.0 L of a solution marked 1.0 M mercury(I) nitrate?
What is the mass of fish, in kilograms, that one would have to consume to obtain a fatal dose of mercury, if the fish contains 30 parts per million of mercury by weight? (Assume that all the mercury from the fish ends up as mercury(II) chloride in the body and that a fatal dose is 0.20 g of HgCl2.) How many pounds of fish is this?
The elements sodium, aluminum, and chlorine are in the same period.
(a) Which has the greatest electronegativity?
(b) Which of the atoms is smallest?
(c) Write the Lewis structure for the simplest covalent compound that can form between aluminum and chlorine.
(d) Will the oxide of each element be acidic, basic, or amphoteric?
What is tin pest, also known as tin disease?
Is the reaction of rubidium with water more or less vigorous than that of sodium? How does the rate of reaction of magnesium compare?
18.2 Occurrence and Preparation of the Representative Metals
Write an equation for the reduction of cesium chloride by elemental calcium at high temperature.
Why is it necessary to keep the chlorine and sodium, resulting from the electrolysis of sodium chloride, separate during the production of sodium metal?
Give balanced equations for the overall reaction in the electrolysis of molten lithium chloride and for the reactions occurring at the electrodes. You may wish to review the chapter on electrochemistry for relevant examples.
The electrolysis of molten sodium chloride or of aqueous sodium chloride produces chlorine.
Calculate the mass of chlorine produced from 3.00 kg sodium chloride in each case. You may wish to review the chapter on electrochemistry for relevant examples.
What mass, in grams, of hydrogen gas forms during the complete reaction of 10.01 g of calcium with water?
How many grams of oxygen gas are necessary to react completely with 3.01 × 1021 atoms of magnesium to yield magnesium oxide?
Magnesium is an active metal; it burns in the form of powder, ribbons, and filaments to provide flashes of brilliant light. Why is it possible to use magnesium in construction?
Why is it possible for an active metal like aluminum to be useful as a structural metal?
What is the common ore of tin and how is tin separated from it?
A chemist dissolves a 1.497-g sample of a type of metal (an alloy of Sn, Pb, Sb, and Cu) in nitric acid, and metastannic acid, H2SnO3, is precipitated. She heats the precipitate to drive off the water, which leaves 0.4909 g of tin(IV) oxide. What was the percentage of tin in the original sample?
Consider the production of 100 kg of sodium metal using a current of 50,000 A, assuming a 100% yield.
(a) How long will it take to produce the 100 kg of sodium metal?
(b) What volume of chlorine at 25 °C and 1.00 atm forms?
What mass of magnesium forms when 100,000 A is passed through a MgCl2 melt for 1.00 h if the yield of magnesium is 85% of the theoretical yield?
18.3 Structure and General Properties of the Metalloids
Give the hybridization of the metalloid and the molecular geometry for each of the following compounds or ions. You may wish to review the chapters on chemical bonding and advanced covalent bonding for relevant examples.
(a) GeH4
(b) SbF3
(c) Te(OH)6
(d) H2Te
(e) GeF2
(f) TeCl4
(g) SiF62−
(h) SbCl5
(i) TeF6
Write a Lewis structure for each of the following molecules or ions. You may wish to review the chapter on chemical bonding.
(a) H3BPH3
(b) BF4−
(c) BBr3
(d) B(CH3)3
(e) B(OH)3
Describe the hybridization of boron and the molecular structure about the boron in each of the following:
(a) H3BPH3
(b) BF4−
(c) BBr3
(d) B(CH3)3
(e) B(OH)3
Using only the periodic table, write the complete electron configuration for silicon, including any empty orbitals in the valence shell. You may wish to review the chapter on electronic structure.
Write a Lewis structure for each of the following molecules and ions:
(a) (CH3)3SiH
(b) SiO44−
(c) Si2H6
(d) Si(OH)4
(e) SiF62−
Describe the hybridization of silicon and the molecular structure of the following molecules and ions:
(a) (CH3)3SiH
(b) SiO44−
(c) Si2H6
(d) Si(OH)4
(e) SiF62−
Describe the hybridization and the bonding of a silicon atom in elemental silicon.
Classify each of the following molecules as polar or nonpolar. You may wish to review the chapter on chemical bonding.
(a) SiH4
(b) Si2H6
(c) SiCl3H
(d) SiF4
(e) SiCl2F2
Silicon reacts with sulfur at elevated temperatures. If 0.0923 g of silicon reacts with sulfur to give 0.3030 g of silicon sulfide, determine the empirical formula of silicon sulfide.
Write a balanced equation for the reaction of elemental boron with each of the following (most of these reactions require high temperature):
(a) F2
(b) O2
(c) S
(d) Se
(e) Br2
Write a formula for each of the following compounds:
(a) silicon dioxide
(b) silicon tetraiodide
(c) silane
(d) silicon carbide
(e) magnesium silicide
From the data given in Appendix I , determine the standard enthalpy change and the standard free energy change for each of the following reactions:
(a) BF3(𝑔)+3H2O(𝑙)⟶B(OH)3(𝑠)+3HF(𝑔)
(b) BCl3(𝑔)+3H2O(𝑙)⟶B(OH)3(𝑠)+3HCl(𝑔)
(c) B2H6(𝑔)+6H2O(𝑙)⟶2B(OH)3(𝑠)+6H2(𝑔)
A hydride of silicon prepared by the reaction of Mg2Si with acid exerted a pressure of 306 torr at 26 °C in a bulb with a volume of 57.0 mL. If the mass of the hydride was 0.0861 g, what is its molecular mass? What is the molecular formula for the hydride?
Suppose you discovered a diamond completely encased in a silicate rock. How would you chemically free the diamond without harming it?
18.4 Structure and General Properties of the Nonmetals
Carbon forms a number of allotropes, two of which are graphite and diamond. Silicon has a diamond structure. Why is there no allotrope of silicon with a graphite structure?
Nitrogen in the atmosphere exists as very stable diatomic molecules. Why does phosphorus form less stable P4 molecules instead of P2 molecules?
Write balanced chemical equations for the reaction of the following acid anhydrides with water:
(a) SO3
(b) N2O3
(c) Cl2O7
(d) P4O10
(e) NO2
Determine the oxidation number of each element in each of the following compounds:
(a) HCN
(b) OF2
(c) AsCl3
Determine the oxidation state of sulfur in each of the following:
(a) SO3
(b) SO2
(c) SO32−
Why does white phosphorus consist of tetrahedral P4 molecules while nitrogen consists of diatomic N2 molecules?
18.5 Occurrence, Preparation, and Compounds of Hydrogen
The reaction of calcium hydride, CaH2, with water can be characterized as a Lewis acid-base reaction:
CaH2(𝑠)+2H2O(𝑙)⟶Ca(OH)2(𝑎𝑞)+2H2(𝑔)
Identify the Lewis acid and the Lewis base among the reactants. The reaction is also an oxidation-reduction reaction. Identify the oxidizing agent, the reducing agent, and the changes in oxidation number that occur in the reaction.
In drawing Lewis structures, we learn that a hydrogen atom forms only one bond in a covalent compound. Why?
What mass of CaH2 is necessary to react with water to provide enough hydrogen gas to fill a balloon at 20 °C and 0.8 atm pressure with a volume of 4.5 L? The balanced equation is:
CaH2(𝑠)+2H2O(𝑙)⟶Ca(OH)2(𝑎𝑞)+2H2(𝑔)
18.6 Occurrence, Preparation, and Properties of Carbonates
Carbon forms the CO32− ion, yet silicon does not form an analogous SiO32− ion. Why?
Complete and balance the following chemical equations:
(a) hardening of plaster containing slaked lime
Ca(OH)2+CO2⟶
(b) removal of sulfur dioxide from the flue gas of power plants
CaO+SO2⟶
(c) the reaction of baking powder that produces carbon dioxide gas and causes bread to rise
NaHCO3+NaH2PO4⟶
Heating a sample of Na2CO3⋅xH2O weighing 4.640 g until the removal of the water of hydration leaves 1.720 g of anhydrous Na2CO3. What is the formula of the hydrated compound?
18.7 Occurrence, Preparation, and Properties of Nitrogen
Write the Lewis structures for each of the following:
(a) NH2−
(b) N2F4
(c) NH2−
(d) NF3
(e) N3−
For each of the following, indicate the hybridization of the nitrogen atom (for N3−, the central nitrogen).
(a) N2F4
(b) NH2−
(c) NF3
(d) N3−
Determine the oxidation state of nitrogen in each of the following. You may wish to review the chapter on chemical bonding for relevant examples.
(a) NCl3
(b) ClNO
(c) N2O5
(d) N2O3
(e) NO2−
(f) N2O4
(g) N2O
(h) NO3−
(i) HNO2
(j) HNO3
For each of the following, draw the Lewis structure, predict the ONO bond angle, and give the hybridization of the nitrogen. You may wish to review the chapters on chemical bonding and advanced theories of covalent bonding for relevant examples.
(a) NO2
(b) NO2−
(c) NO2+
How many grams of gaseous ammonia will the reaction of 3.0 g hydrogen gas and 3.0 g of nitrogen gas produce?
Although PF5 and AsF5 are stable, nitrogen does not form NF5 molecules. Explain this difference among members of the same group.
The equivalence point for the titration of a 25.00-mL sample of CsOH solution with 0.1062 M HNO3 is at 35.27 mL. What is the concentration of the CsOH solution?
18.8 Occurrence, Preparation, and Properties of Phosphorus
Write the Lewis structure for each of the following. You may wish to review the chapter on chemical bonding and molecular geometry.
(a) PH3
(b) PH4+
(c) P2H4
(d) PO43−
(e) PF5
Describe the molecular structure of each of the following molecules or ions listed. You may wish to review the chapter on chemical bonding and molecular geometry.
(a) PH3
(b) PH4+
(c) P2H4
(d) PO43−
Complete and balance each of the following chemical equations. (In some cases, there may be more than one correct answer.)
(a) P4+Al⟶
(b) P4+Na⟶
(c) P4+F2⟶
(d) P4+Cl2⟶
(e) P4+O2⟶
(f) P4O6+O2⟶
Describe the hybridization of phosphorus in each of the following compounds: P4O10, P4O6, PH4I (an ionic compound), PBr3, H3PO4, H3PO3, PH3, and P2H4. You may wish to review the chapter on advanced theories of covalent bonding.
What volume of 0.200 M NaOH is necessary to neutralize the solution produced by dissolving 2.00 g of PCl3 is an excess of water? Note that when H3PO3 is titrated under these conditions, only one proton of the acid molecule reacts.
How much POCl3 can form from 25.0 g of PCl5 and the appropriate amount of H2O?
Write equations showing the stepwise ionization of phosphorous acid.
Draw the Lewis structures and describe the geometry for the following:
(a) PF4+
(b) PF5
(c) PF6−
(d) POF3
Why does phosphorous acid form only two series of salts, even though the molecule contains three hydrogen atoms?
Assign an oxidation state to phosphorus in each of the following:
(a) NaH2PO3
(b) PF5
(c) P4O6
(d) K3PO4
(e) Na3P
(f) Na4P2O7
Phosphoric acid, one of the acids used in some cola drinks, is produced by the reaction of phosphorus(V) oxide, an acidic oxide, with water. Phosphorus(V) oxide is prepared by the combustion of phosphorus.
(a) Write the empirical formula of phosphorus(V) oxide.
(b) What is the molecular formula of phosphorus(V) oxide if the molar mass is about 280.
(c) Write balanced equations for the production of phosphorus(V) oxide and phosphoric acid.
(d) Determine the mass of phosphorus required to make 1.00 × 104 kg of phosphoric acid, assuming a yield of 98.85%.
18.9 Occurrence, Preparation, and Compounds of Oxygen
Using equations, describe the reaction of water with potassium and with potassium oxide.
Write balanced chemical equations for the following reactions:
(a) zinc metal heated in a stream of oxygen gas
(b) zinc carbonate heated until loss of mass stops
(c) zinc carbonate added to a solution of acetic acid, CH3CO2H
(d) zinc added to a solution of hydrobromic acid
Write balanced chemical equations for the following reactions:
(a) cadmium burned in air
(b) elemental cadmium added to a solution of hydrochloric acid
(c) cadmium hydroxide added to a solution of acetic acid, CH3CO2H
Write balanced chemical equations for the following reactions:
(a) metallic aluminum burned in air
(b) elemental aluminum heated in an atmosphere of chlorine
(c) aluminum heated in hydrogen bromide gas
(d) aluminum hydroxide added to a solution of nitric acid
Write balanced chemical equations for the following reactions:
(a) sodium oxide added to water
(b) cesium carbonate added to an excess of an aqueous solution of HF
(c) aluminum oxide added to an aqueous solution of HClO4
(d) a solution of sodium carbonate added to solution of barium nitrate
(e) titanium metal produced from the reaction of titanium tetrachloride with elemental sodium
What volume of 0.250 M H2SO4 solution is required to neutralize a solution that contains 5.00 g of CaCO3?
Write a balanced chemical equation for the reaction of an excess of oxygen with each of the following. Remember that oxygen is a strong oxidizing agent and tends to oxidize an element to its maximum oxidation state.
(a) Mg
(b) Rb
(c) Ga
(d) C2H2
(e) CO
Which is the stronger acid, H2SO4 or H2SeO4? Why? You may wish to review the chapter on acid-base equilibria.
18.10 Occurrence, Preparation, and Properties of Sulfur
Explain why hydrogen sulfide is a gas at room temperature, whereas water, which has a lower molecular mass, is a liquid.
Which is the stronger acid, NaHSO3 or NaHSO4?
Which is a stronger acid, sulfurous acid or sulfuric acid? Why?
Give the Lewis structure of each of the following:
(a) SF4
(b) K2SO4
(c) SO2Cl2
(d) H2SO3
(e) SO3
Explain why sulfuric acid, H2SO4, which is a covalent molecule, dissolves in water and produces a solution that contains ions.
18.11 Occurrence, Preparation, and Properties of Halogens
What does it mean to say that mercury(II) halides are weak electrolytes?
The following reactions are all similar to those of the industrial chemicals. Complete and balance the equations for these reactions:
(a) reaction of a weak base and a strong acid
NH3+HClO4⟶
(b) preparation of a soluble silver salt for silver plating
Ag2CO3+HNO3⟶
(c) preparation of strontium hydroxide by electrolysis of a solution of strontium chloride
SrCl2(𝑎𝑞)+H2O(𝑙)−→−−−−−−−electrolysis
What is the hybridization of iodine in IF3 and IF5?
Predict the molecular geometries and draw Lewis structures for each of the following. You may wish to review the chapter on chemical bonding and molecular geometry.
(a) IF5
(b) I3−
(c) PCl5
(d) SeF4
(e) ClF3
Which halogen has the highest ionization energy? Is this what you would predict based on what you have learned about periodic properties?
Explain why, at room temperature, fluorine and chlorine are gases, bromine is a liquid, and iodine is a solid.
What is the oxidation state of the halogen in each of the following?
(a) H5IO6
(b) IO4−
(c) ClO2
(d) ICl3
(e) F2
Physiological saline concentration—that is, the sodium chloride concentration in our bodies—is approximately 0.16 M. A saline solution for contact lenses is prepared to match the physiological concentration. If you purchase 25 mL of contact lens saline solution, how many grams of sodium chloride have you bought?
18.12 Occurrence, Preparation, and Properties of the Noble Gases
Give the hybridization of xenon in each of the following. You may wish to review the chapter on the advanced theories of covalent bonding.
(a) XeF2
(b) XeF4
(c) XeO3
(d) XeO4
(e) XeOF4
What is the molecular structure of each of the following molecules? You may wish to review the chapter on chemical bonding and molecular geometry.
(a) XeF2
(b) XeF4
(c) XeO3
(d) XeO4
(e) XeOF4
Indicate whether each of the following molecules is polar or nonpolar. You may wish to review the chapter on chemical bonding and molecular geometry.
(a) XeF2
(b) XeF4
(c) XeO3
(d) XeO4
(e) XeOF4
What is the oxidation state of the noble gas in each of the following? You may wish to review the chapter on chemical bonding and molecular geometry.
(a) XeO2F2
(b) KrF2
(c) XeF3+
(d) XeO64−
(e) XeO3
A mixture of xenon and fluorine was heated. A sample of the white solid that formed reacted with hydrogen to yield 81 mL of xenon (at STP) and hydrogen fluoride, which was collected in water, giving a solution of hydrofluoric acid. The hydrofluoric acid solution was titrated, and 68.43 mL of 0.3172 M sodium hydroxide was required to reach the equivalence point. Determine the empirical formula for the white solid and write balanced chemical equations for the reactions involving xenon.
Basic solutions of Na4XeO6 are powerful oxidants. What mass of Mn(NO3)2•6H2O reacts with 125.0 mL of a 0.1717 M basic solution of Na4XeO6 that contains an excess of sodium hydroxide if the products include Xe and solution of sodium permanganate?