21-1 Periodic Trends and Charge Density

• The group 1 atom in a given period is the largest that hydrogen is often placed and is the most easily ionized.
• The group 1 elements have a low densi table, but some of them are not ties and are an alkali metal.
• This metal is highly reactive.
• The reactivity of the alkali metals can be seen in their reactions with water.

• The group 2 has been measured.

• The metals of group 2 are not as dense as a typical metal, but they are still highly reactive and less dense than the alkali metals.

• The group 2 metals have densities that are greater than that of water, and they only react slowly with water.

• Group 13 and group 14 will be discussed in this chapter.
• Both metals and nonmetals are encountered in these groups.
• Boron has interesting chemistry because it tends to form molecule with incomplete octets around the central boron atoms.
• One of the most widely used metals is aluminum.
• It is possible to get aluminum metal from its compounds.
• Because aluminum production requires a lot of electricity, aluminum-production plants are located close to a lot of hydroelectricity.

• Group 13--Gallium, indium, and thallium--are all metals.
• The chemistry of group 13 is dominated by aluminum and boron, and we will only mention the heavier elements in this chapter.
• Group 14 has a nonmetal, two metalloids, and two metals.
• The chemistry of carbon is the most important in the group since it occurs in all living systems.
• Silicon is found in many minerals.
• Tin and lead can be obtained using methods that have been used for thousands of years.

• There are many opportunities to relate new information to principles presented earlier in the text.
• The ideas of atomic structure, periodic trends in atomic and ionic radii, chemical bonding, and thermodynamics will help us understand the chemical behavior of the elements.

• The periodic trends that we have covered in this text can be rationalized in the chemistry of the elements.
• The elements in a given group have similar electronic configurations, but not the same chemical properties, because the atoms of each group have similar electronic configurations.
• The lightest member of a group has features that are different from the rest of the group.

• Trends in atomic properties will be reviewed in this section.
• We will begin to understand the trends in the chemistry of the elements with these ideas.

• The atomic properties of an element are responsible for its chemistry.
• The electronegativity of an atom is important.

• The ground state electronic configuration of He is 1s2.

• A summary of trends in atomic radius, first ionization energy, electron affinity, electronegativity, and atomic polarizability can be found here.

• The shaded elements are the focus of the chapter.

• Chapter 9 discussed atomic radii, ionising energies, and electron affinities.

• Chapter 10 and Chapter 12 discussed electronegativities and polarizabilities.

• We have discussed polarization of the anion trends before in this text.

• A high charge density distorts the electron from top to bottom in a group.
• The anion has first ion energies, electron affinities, and cloud around it.
• The electronegativities show that the quantities increase across the electron cloud and decrease down a group.
• cations are smaller than the parent atoms when shown with a dashed line and radii, and anions are larger, so it is important to remember.
• Explanations for the trends have been shown.
• Chapters 9 and 10 have been given elsewhere in the text.

• The bond between the anion electron cloud and the internuclear region is distorted when a cation interacts with an anion.

• As a result of the distortion, the bond between the cation and the anion has a variety of different character.

• Charge density is deterred by the polarizability of the anion and the cation.
• The electron cloud of the anion is distorted for me.
• The larger and more polarizable anions derived from atoms that are lower down in a group are defined by some authors as charge density.
• The I- ion is more polarizable than the F-.

• The charge density of the Main- Group Elements I: Groups 1, 2, 13, and 14 is between 1 and 1000 Cmml -3.

• The charge density increases as the charge on the cation increases.

• The higher the charge density, the greater the ability of a cation to distort the electron cloud of an anion toward itself.

• The charge density concept will be used to rationalize certain observations.
• It will be used to help us understand why there are dramatic differences in the properties of elements in the same group.
• It is not possible to use a single quantity of Group 1 Elements as a substitute for careful consideration of all contributing factors.

• 3 and AlI3 are used to measure the amount of solid crust.

• Since prehistoric times, some of their compounds have been used.
• The elements were isolated in pure form about 200 years ago.
• The elements of the alkali metals are difficult to break down by ordinary chemical means, so discovery had to wait for new scientific developments.
• The two substances were discovered through electrolysis.
• It was discovered in 1817.
• Cesium and rubidium were identified as new elements.
• Francium was isolated from actinium.

• Natural brines can be used to obtain a number of Li, Na, and K compounds.
• Na Albert Russ/Shutterstock 2CO3 can be mined as a solid deposit.
• The salt is obtained from the water.
• LiAl1SiO322 is the name for Rubidium and cesium.

• The group 1 elements are the most active metals.
• Several of their properties are listed in Table 21.2, and a few of them are discussed next.

• ion pairs are converted to gaseous atoms when NaCl is vaporized.
• As excited atoms 1Na*2 return to their groundstate electron configurations, light with a wavelength of 589 nm emits as the excited atoms Na(g) are excited to higher energies.

• Yellow2 Alkali metal compounds are used in fireworks.

• The atomic radii of the group 1 elements increase from the top to the bottom, as was described in Chapter 9.
• These large atoms make Chip Clark/Fundamental Photographs for a relatively low mass per unit volume.
• The lighters of the alkali metals will float.

• The property leads to soft met The sodium, an active metal with low melting points.
• A bar of sodium is covered with a thick oxide coating.

• The values given here assume a coordination number of 4 for Li+ and 6 for the others.

• Ten minerals are ranked on the Mohs scale, ranging from that of talc to diamond.
• Only substances with lower values can be scratched.

• A good indicator of the extreme metallic character of the group 1 elements is their standard reduction potentials, which are large, negative quantities.
• The metal M(s) is very easy to oxidize to M+1aq2 because it is difficult to reduce the ion M+.
• The alkali metals can reduce water to H21g2.

• We can use the values given in Table 21.2 to calculate them.

• The Edegcell values show that the strongest reducing agent in the solution is lithium.
• The strong reducing agents in the alkali metals are in the solution.
• The equilibrium position for reac is always very far to the right, and so the reaction is controlled by completion regardless of which alkali metal is involved.

• The time it takes for a reaction to happen and the rate of it.
• It is necessary to consider what happens to the energy that is released by the controlled factors to explain this observation.

• The energy released by the reaction is used to heat the system.
• The energy released by the reaction is enough to melt the unreacted metal.

• The reaction of lithium and water is not as vigorous as it is for the other alkali metals because it does not melt as the reaction proceeds.

• The melting point of NaCl is too high a temperature to carry this economically.

• The Downs cell is used for the reduction of molten KCl.

• The Downs cell has NaCl1l2 and K1g2 in it.
• At low temperatures, most of the KCl(l) remains is molten NaCl(l) to which CaCl2 has been added unreacted.
• The equilibrium is displaced far to the right as to lower the melting point of K(g) from the molten mixture.
• The K(g) is free of any Na(g) present when the liquid metals are fractionally distilled.
• Rb and Cs can be produced with Ca metal as the reducing agent.

• The most important use of the metal is because it is so easy to oxidize and because it can be kept apart by reducing agent.

• Titanium metal can be obtained from the reduction of TiCl4 by Na.

• In a nuclear reactor, sodium is used as a heat-transfer medium.

• It has better thermal con elevated temperature.

• It is easy to pump because of its low density and low viscosity.
• sodium is used in outdoor lighting.
• The total quantity consumed in this application is small because each lamp uses a small amount of Na.

• It is possible to make high-strength, low density alloys with aluminum and magnesium using lithium metal.
• These are used in the aircraft industry.
• The ease of oxidation and the large number of electrons produced by a small mass of lithium make it an anode material in batteries.

• It takes 6.94 g Li to produce one mole of electrons.
• In cardiac pacemakers, the installed battery must have high reliability and a long lifetime to be useful.
• There is an X-ray photograph of a pacemaker.

• Li is the smallest of the alkali metal atoms.

• It is a different matter when it comes to oxidation.
• The result of a hypothetical three-step process is what we can think of.

• We must compare tendencies in each of the three steps if we want to form M+1aq2 by oxidation of the metals.

• The secondary hydration sphere is formed by the listing of electrode potentials holding other molecule, but more weakly.

• The easiest substance to oxidize is Li.

• A common method of summing up important reactions is from a compound of central importance.
• This section describes some of the reactions.
• Alternative methods may be used to prepare a number of these compounds.
• The conversion of Na2CO3 to NaOH is no longer important.

• Group 1 metals, Li-Fr, have a valence configuration of ns1 and are only found in the oxidation state.
• Most of the compounds of group 1 metals are stable.
• A large amount of information is difficult to organize when studying the chemistry of the elements.

• Similar diagrams can be constructed for other compounds, even though this diagram deals with sodium compounds.

• Some of the conversions occur in one step, such as the reaction of NaCl with H2SO4 to form Na2SO4.
• The preparation of Na2CO3 involves two or more reactions.
• The principal reactants are written near the arrows.
• A C/ symbol is used to indicate if a reaction mixture must be heated.
• The by-products may not be noted in the diagram.

• The cations are hydrated when salts are dissolved in water.
• The anions have the same hydration but with slightly positive hydrogen atoms in the water.
• Water is a part of the solid structure when a salt is crystallized.
• A number of factors must be considered before a simple rule can be used for predicting whether the ion will retain their hydration spheres in the solid state.
• cations with high charge densities tend to retain all or part of their hydration spheres in the solid state.
• When the cations have low charge densities, they lose their hydration spheres and form anhydrous salts.

• The charge densities of the alkali metals are shown in Table 21.2, but the majority of the salts are anhydrous.
• Salts are most likely to be hydrated.

• Figure 21-5 shows a route from NaCl to Na2CO3.
• The other necessary reactants are noted for each conversion.
• Keep in mind that other reaction products must be included in balanced chemical equations for each conversion.

• Na2SO41s2 is produced from concentrated sulfuric acid.

• The Na2SO4 is reduced to Na2S with 4 C1s2 C/ and 4 CO1g2 carbon.

• The reaction between Na2S and CaCO3 is the final step.

• There are three separate chemical equations.
• One of the sodium compounds on the route chosen is produced by a balanced chemical equation.

• Write chemical equations for the reactions that take place.

• Write chemical equations for the reactions that take place.

• The most important of the ionic halides are NaCl and KCl.

• Salt is used in the production of chemicals.
• It is not listed among the top chemicals because it is a raw material.
• Large quantities of NaCl can be obtained through the use of seawater.

• Salt is used to preserve meat and fish.
• In the chemical industry, NaCl is a source of many chemicals.

• There are naturally occurring brines that give rise to KCl.
• It is used most in plant fertilization because of its importance to plant growth.
• Sea salt stacks that have been Harvested by Sea Salt Stacks that have been Harvested by Sea Salt Stacks that have been Harvested by Sea Salt Stacks that have been Harvested by Sea Salt Stacks that have been Harvested by Sea Salt Stacks that have been Harvested

• The Main- Group Elements I: Groups 1, 2, 13, and 14 chloride are used as a raw material in the manufacture of KOH, KNO3 and other industrially important potassium compounds.

• The hydride ion, H-, and the sodium chloride structure are found in alkali metal hydride.
• The alkali metal hydride is very reactive.

• The metal hydride reacts with metal halides.
• An example of this is the reaction of LiAlH4 with aluminum chloride.
• LAH is a reducing agent used in chemistry.

• Adding finely divided LiH to a solution of AlCl3 in a nonaqueous solvent is how the reaction is carried out.
• Both LiAlH4 and LiH react vigorously with water, so a nonaque ous solvent is used.

• LiAlH4 is a white solid because of careful and controlled evaporation of the solvent.

• The alkali metals react quickly with oxygen to produce oxides and Hydroxides.
• The oxide M2O can be prepared if the supply of oxygen is carefully controlled.
• The Na2O2 and Na2O react with excess oxygen.
• Ionic compounds include the oxides, peroxides, and superoxides.

• The table shows the principal products of the reactions of the alkali metals.

• RbO2 is known.

• As we move down the group from Li to Cs, we notice that the oxide (M2O) shifts to the superoxide (MO2).

• The plots for the alkali metals are from M(s) and O21g2.

• The M(s) and O21g2 are not the same as the peroxides and monoxides.

• These reactions are almost the same as the O2 M2O.

• The stability can be measured relative to some set of refer not dissipated immediately.
• The O2 is not very stable compared to oxygen.
• The process temperature for the formation of the O2 is very warm.

• The O atom is called Li2O1s2.

• 2 favored is O1g2 C/rH.

• It is surprising that M2O1s2 is not formed because of the large energy requirement.

• The lattice energy is the enthalpy change for the process.

• The formation of Li2O1s2 starts from Li1s2 and O21s2 because the lattice energy of Li2O is very negative.

• The formation of the M2O lattice is less favorable due to the smaller lattice energies of Na2O, K2O, Rb2O, and Cs2O.

• The heavier alkali metals react with excess oxygen to give either M2O2 or MO2.

• Li2O21s2 is a minor product, and this suggests that it is a stable product.

• The other alkali metal peroxides have to be heated to higher temperatures.

• A molecule is formed when the oxygen atom combines with another oxygen atom.

• There are many uses for the peroxides.
• It is possible to use a powerful oxidant and a bleaching agent.
• It is possible to use the superoxide,KO2, for this purpose.
• Basic solutions can be formed by the reaction of the oxides, peroxides, and superoxides of the alkali metals.
• The acid-base reaction that produces the alkali metal hydroxide is the reaction of an alkali metal oxide with water.

• One mole of oxide produces two moles of hydroxide ion.

• The water reacts with the peroxide ion in a similar way to produce hydrox ide ion and hydrogen peroxide.

• 2 OH-1aq2 + H2O21aq2 Hydrogen peroxide slowly disproportionates into water and oxygen.

• The water reacts with the superoxide ion to give it hydrogen, hydrogen per oxide, and oxygen.

• The strong bases of the group 1 metals are due to the dissociation of the hydroxides.
• In Section 19-8, we learned about the commercial production of sodium hydroxide.

• H2O is reduced to H21g2 and Na+1aq2 goes through the electrolysis unchanged.
• Both of the hydroxides are made in a similar fashion.
• The reaction of the group 1 metals with water can be used to make alkali metal hydroxides.

• The manufacture of soaps and detergents involves the use of alkali hydroxides.

• The M2CO3 can be heated to a high temperature of 17800 degreesC2 and can be killed by decomposing to M2O1s2 and CO21g2.
• The chemistry of the second member of the group is usually representative of the group.
• The first member is not a representative of the group.

• The treatment seems to affect the balance of Na+ and that of K+ and Ca+ in the cell.

• The earliest 2CO3 NaHCO3 n H2O was found in dry lakes in California.

• A simplified description of the process is given by the following equation.

• The reaction is carried out in two stages.
• In the United States, ammonia is bubbled into a concentrated brine solution.

• Solid arrows trace the main reaction sequence.
• dashed arrows show recycling reactions.

• CO2 is bubbled through the ammoniated brine in the chemistry of the Main- Group Elements I.
• The reaction shows that the sodium bicarbonate can be isolated and sold or converted to sodium carbonate by heating.

• The Solvay process only involves simple precipitation and acid-base reactions.
• In one step, materials are recycled into another step.
• A process that recycles materials reduces the use of raw materials and cuts down on the production of by-products, which are an expense in disposal.
• CaO is also used when limestone 1CaCO32 is heated to produce the reactant CO2.
• It is converted to Ca1OH22, which is used to convert NH4Cl to NH31g2.
• The NH31g2 is used in the production of ammoniated brine.

• The Solvay process only has one by-product which is limited in demand.
• In the past, some CaCl2 was used for deic ing roads in the winter and for dust control on dirt roads in the summer.
• The majority of the CaCl2 was dumped into local lakes and streams.
• Dumping is no longer allowed by environmental regulations.
• Natural sources of sodium carbonate have replaced the Solvay process in the United States due to the regulations.
• Elsewhere in the world, the process is still used.

• The process is based on reactions.

• The paper industry uses Na2SO4 the most.
• In the process of papermaking, undesirable lignin is removed from wood by using an alkaline solution of Na2S.
• The reduction of Na2SO4 with carbon is required for this step.

• 45 kilograms of Na2SO4 is required for every metric ton of paper produced.

• Most of the NH3 complexes have relatively low charge densities.

• In 1967, Charles J. Pedersen reported the discovery of a type of Lewis base.

• The 12-crown-4 ether consists of a ring of 12 atoms, of which 4 are oxygen atoms.

• The structure of crown ethers is reflected in the name it is given.

• Oxygen atoms give electron density to the metal ion.

• One of the factors that affect which cations bind to which crown ether is the Cavity size.

• The values of the equilibrium constants aren't different enough to make 18-crown-6selective for just K+.

• Synthetic organic chemistry exploits the formation of crown ether complexes to form ionic reagents that can be dissolved in nonpolar solvents.
• KMnO4 is not in pure benzene, but 3K118-crown-624 is.

• Soap is an emulsi common soaps, it is used to describe certain synthetic products, such as sodium lauryl sulfate, whose manufacture involves the following formation of an emulsion of conversions.

• The lauryl sulfate anion has a long, non polar tail.
• Structural features are common to detergents and soaps.

• The product of the reaction of palmitic acid and NaOH is called sodium palmitate.

• Notice that the anion of the soap has a long, nonpolar tail and a polar head.

• The soaps that are hard are called sodium soaps.
• Soft soaps have low melting points.
• A soap called 1C17H35CO2Li2 is used to make greases.
• At both high and low temperatures, these greases have lubricating and water-repellent properties.
• Under the conditions in which oil would run off, greases remain in contact with metal parts.

• There is a long nonpolar portion buried in a droplet of oil and a polar head projecting into an aqueous medium.

• There is a reason for this 2O.

• Balance the equations for the two reactions.

• The elements of group 1 and group 2 have the same alkaline earth metals.

• The term is crust.

• The heavier group 2 metals-- Ca, Sr, Ba, and Ra-- are more active in chemistry than the group 1 metals.
• In terms of certain physical properties, the group 2 elements are more metallic than the group 1 elements, as we can see by comparing Tables 21.2 and 21.4.

• The table shows that beryllium is out of step with the other group.
• It is harder than the others and has a higher melting point.
• Its chemical properties are different.
• BeO does not react with water.

• The coordination number is 6 and the aIonic radii are green.

• Poor conductors of electricity, BeF2 and BeCl2 are in the molten state.

• The high charge density of the beryllium cation is related to the unusual chemical behavior of beryllium.
• The small Be2+ ion polarizes any nearby anion, drawing electron density toward itself, creating a bond with significant covalent character.
• Some of the compounds of beryllium have some characteristics that are similar to covalent solids.

• electron density is donated from H2O or OH- to Be2+ because of the high polarizing power of the Be2+ ion and the resulting com Tetrahedral shape of the plex has a well-defined structure.
• The 3Be1OH22442+ ion is shown in the figure in the margin.

• The ion 3Be1H2O2442+ has a tetrahedral structure.

• Be are bonds to one Be atom.

• The bonds are made with lone-pair electrons on the atoms.

• Two types of bonds are indistinguishable.

• Reducing their salts with other active metals is the preferred method of producing the group 2 metals.
• The source of beryllium compounds is Be3Al2Si6O18.
• This mineral is reduced to give BeF2 after being processed.
• When low density is a primary requirement, beryllium metal is used.
• Be is used in springs, clips, and electrical contacts because of its resistance to metal fatigue.
• The Be atom does not absorb X-rays or neutrons, so beryllium is used to make windows for X-ray tubes.
• Because they are toxic, beryllium and its compounds are limited in their use.
• They are suspected of being carcinogens even at a low level.

• The reduction of their oxides with aluminum can be used to obtain calcium, strontium, and barium.
• Calcium metal is used as a reducing agent to prepare from oxides and other metals.
• Some of the compounds of Strontium and barium are important.
• There are some salts of Ba that provide vivid colors.

• The electrolyte is made of molten Na, Ca, and Mg.

• The process for making NaHCO3 is similar to the Solvay process.

• Natural brines are the source of magnesium.
• The abundance of this material in the water is 1350.
• The precipitation of Mg1OH221s2 with slaked lime is the first step in the process.

• Slaked lime is formed when quicklime reacts with water.

• The precipitated Mg1OH221s2 is washed and dissolved.

• Pure Mg metal and Cl21g2 can be found in the dried concentration.
• The Cl21g2 is recycled.

• Magnesium has a lower density than other metals.
• magnesium is used in the production of lightweight objects, such as aircraft parts.
• Magnesium is used in a number of processes, including the production of beryllium.
• The ease with which magnesium is oxidation underlies its use in sacrificial anodes.
• Magnesium can be used in firework as it burns in the air with a white light.

• The magnesium burns in an atmosphere of carbon dioxide, which shows the good reducing properties of the metal.
• Write an equation for this reaction.

• The +2 oxidation state is where the alkaline earth metals are found.
• The ns2 electrons in the group 2 metals are lost when they combine with nonmetals to form compounds.
• The alkaline earth metals form mostly ionic compounds, but covalent bonding is visible in magnesium compounds and especially in beryllium compounds.

• Group 2 compounds have different properties than group 1 com pounds.
• The larger ionic charge of group 2 cations may be the reason for the difference.

• This difference in lattice energy helps explain why NaOH is verysoluble in water up to about 20 M NaOH(aq).
• The heavier group 2 hydroxides are more stable.
• The carbonates, fluorides, and oxides are alkaline earth compounds.

• The standard method for preparing MX21s2 in anhydrous form is to dehydrate the hydrates obtained from the metal and hydrohalic acid reaction.
• The method of preparation cannot be used to prepare beryllium halides.

• Anhydrous BeCl2 is prepared from BeO and CCl4.

• The halides have different uses.
• The preparation of magnesium metal, fireproofing wood, special cements, ceramics, treating fabrics, and as a refrigeration brine are some of the uses of the substance.

• It is used in water treatment, in the removal of SO21g2 from the smokestack gases in electric power plants, and in the making of Ca1OH22, an important and inexpensive strong base.

• The strong bases of the group 2 metals are the hydroxides.
• It is not very water-soluble, but it is used in a variety of processes.

• The mortar used in bricklaying is composed of a mixture of slaked lime, sand, and water.
• The excess water in the mortar is absorbed by the bricks.

• The mortar is made up of hydrated calcium carbonate and silicate from the sand.

• Reaction is general for all group 2 hydroxides.
• This reaction has been used to preserve art objects.
• A solution of Ba1NO322 is sprayed onto frescos.

• An ammonia solution is applied to the fresco when there are small cracks and spaces.
• Ba1OH22 was formed when the ammonia raised the solution's pH.
• The carbon dioxide from the air reacts with the water.

• The cracking fresco can be strengthened without affecting the delicate colors.

• Suggest an alternative if you disagree with the statement.

• Hydrate is a characteristic of alkaline earth compounds.

• The Ba2+ X is a low charge density and shows little or no tendency to retain its hydration sphere in the solid state.

• Sulfates and carbonates are insoluble in water.
• The most important minerals of the group 2 metals are the compounds.

• Some minerals may be present in limestone.
• The majority of limestone is used as a building stone.
• limestone is used in the manufacture of quicklime and slaked lime, as an ingredient in glass, and as a flux in metallurgical processes.

• It is produced in long kilns like the one shown in the photograph.
• In the kiln, limestone, clay, and sand are heated to higher temperatures as they slowly move down the inclined kiln.
• The water is driven off by the first two things.
• The limestone to CaO1s2 and CO21g2 is then decomposing.
• CaO combines with limestone to form silicates from the sand and clay, which are carried out in a long rotary aluminates.
• Pure cement doesn't have much strength.
• Portland cement is an important material for the construction of bridge CaO and for the manufacture of piers and other underwater structures because of its ability to resist cracking even under Portland cement.

• CaCO3 is used in many products.
• In papermaking, it is used to impart brightness, opacity, smoothness, and good ink-absorbing qualities to paper.
• It is suited to newer papermaking processes that produce acid-free paper with an expected shelf life of 300 years or more.
• CaCO3 is used in a wide range of things, from rubber to food and cosmetics.
• It is used as an antacid and as a supplement for the prevention of osteoporosis, a condition in which the bones become porous and brittle and break easily.

• The limestone reacts at room temperature.
• To prevent the reverse reaction, a high temperature must be used and CO21g2 must be removed from the kiln.

• Limestone 1CaCO32 is responsible for the beautiful natural formations found in limestone caves.
• Natural groundwater is slightly acidic because of dissolved CO21g2 and is a solution of carbonic acid, H2CO3.

• The carbonates are bases and can be dissolved in acidic solutions.
• CaCO3 is converted to Ca1HCO322 when mildly acidic water creeps through limestone beds.

• A limestone cave can be created by dissolving action over time.
• A loss of both water and CO2 and conversion of Ca1HCO3221aq2 back to CaCO31s2 can be caused by the evaporation of the solution.
• This process takes a long time.

• stalactites and stalagmites grow together.

• 50 million metric tons of gypsum are consumed annually in the United States.

• The plaster of Paris reverting to gypsum when mixed with water is CaSO4 2 H2O1s2 + 2 H2O1g2 Because it expands as it sets, a mixture of plaster of Paris and water is useful in making castings where sharp details of an object must be retained.
• In jewelry making and dental work, plaster of Paris is used.
• Producing gypsum wallboard is the most important application in the construction industry.

• The compound BaSO4 Plaster of Paris castings is so insoluble that it is safe to use as a "barium milkshake" to coat the stomach.

• 3Mg1H2O2642+ is the hydrated magnesium ion.

• We can look at these similarities.

• Some of the chemical similarities between magnesium and lithium can be found in the following list.

• The oxide is given by O2 instead of the peroxide.

• The oxide and carbon dioxide can be given from the carbonates of magnesium and lithium.
• The carbonates of the remaining group 1 metals are stable.

• Ge had charge densities.
• The increase in charge density is related to the increase in size of the Mg2+ relative to Li+.

• There is a high degree of covalency in the two elements in each compound.
• Many relationships exist between Be and Al and between B and Si as well.
• The properties of these relationships are similar.

• Write balanced chemical equations for the reactions if you have a reason for it.

• 0.1 moles of M, a group 1 metal, react with enough oxygen to give 0.05 moles of compound X.
• The only product that can be reacted with water is a hydroxide.
• In a separate experiment, 0.1 moles of the metal reacts with water to give 0.1 moles of a hydroxide and 0.05 moles of a gas.
• The metal and compounds are identified.
• Balance the chemical equations for the reactions.

• Some information is provided for each reaction.
• One way to tackle this problem is to write partial chemical equations for the reactions, and then use the information provided to complete the equations.

• M is a group 1 metal, so the hydroxide has the formula MOH.

• Water reacts with the alkali metals to give MOH and H2.
• Y must be H2.
• X could be M2O, M2O2, or MO2 in the first reaction.
• There is only one product in the second reaction.
• X cannot be M2O2 or MO2 because they react with water.
• X must be M2O.
• Li is the only alkali metal that reacts with oxygen to give M2O.

• The solution to this problem was to know that the normal oxide, M2O, reacts with water to give MOH, and that the only alkali metal that gives M2O is lithium.

• The compound X and nitrogen gas are given by the reaction of sodium nitrite and sodium metal.
• The compound X reacts with oxygen to give the compound Y.
• Write balanced chemical equations for the reactions described after identifying X and Y.

• A group 2 metal is heated with carbon at 1100 degC to produce a single compound X.
• Solid carbon is produced along with compound Y when compound X is heated in excess of N21g2.
• The plaster of Paris is made from the group 2 metal sulfate and compound Y.
• Write balanced chemical equations for the reactions described, and describe the shape of the anion in Y.

• Boron has charge and chemical properties.
• The remaining members of group 13-Al, Densities of Group 13 Ga, In, and Tl are metals and will be discussed later in this section.

• For the first time elements in the 3 13 have more than one oxidation Oxidation State state.
• The elements of this group exhibit both oxidation states.
• Boron is a nonmetal.
• covalent bonds are formed by the other Charge Density, members of the group.

• Lewis acids are made strong by this deficiency.
• The bonding of a type that we have not previously encountered is caused by the electron deficiency of some boron compounds.

• The bonding takes place in the boron hydride.

• The molecule BH3 (borane) may exist as a reaction intermedi ate, but it has not been isolated as a stable compound.

• It has six electrons in its valence shell.
• Diborane, B2H6, is the simplest boron hydride that has been isolated.

• Bonding theories fail for this molecule.

• The two B atoms and four of the H atoms are in the same plane as the page.
• Four bonds have eight electrons involved.
• Four electrons are left to bond the two remaining H atoms to the two B atoms and also to bond the B atoms together.

• We have not had a lot of Bonding, but atom bridges are fairly common.

• molecu lar orbital theory can be used to rationalize the bonding in these three-center bonds.
• The structure of B5H9 is described in Figure 21-18.

• New and exciting developments in chemistry are provided by them.

• In Italy, Russia, Tibet, Turkey, and the desert regions of California, concentrated ores can be found.

• There are many color-safe bleaches that use the sodium perborate as a bleach alterna.
• 2 bonds the boron atoms.
• There are five H atoms.

• One of the key compounds from which other boron compounds can be used.
• Bridge pairs of B atoms come from the weakly acidic nature of boric acid.

• 2B4O7 10 H2O is converted to B1OH23 by reaction with H2SO4 B1OH23 is converted to B2O3 when heated strongly.
• Boron can be prepared from B2O3.

• The water is the source of the H3O.

• The perborate ion is used in cleaning agents.
• Boric acid can be used to kill roaches and as an antiseptic in eyewash solutions.
• Boron compounds are used in a wide range of products.

• The Lewis acid behavior of boron compounds can be seen in the halides.

• A coordinate covalent bond is formed between a Lewis acid and a Lewis base, with the pair of electrons coming from the Lewis base.
• The red arrow shows the transfer of electron density from the Lewis base to the Lewis acid.

• F from sp2 to sp3.

• The 2p orbitals on the B 3 molecule may be the reason for the observations.

• There is a greater tendency to form adducts incl 3 than incl 3.

• In the solid phase, the molecule adopts a planar geometry, but in the gas phase it is nonplanar.

• B is usually described in terms of two-electron bonds.

• Key substances are listed in Figure 21-19.
• We can identify other plausible reactants and products by writing an incomplete chemical equation for each reaction.

• H2SO4 is required for the conversion of a salt to a acid.
• This is an acid-base reaction.
• Na2SO4 and H2O are two products of the reaction.
• The equation can be balanced by inspection because the reaction does not involve changes in oxidation states.
• B2O3 can be converted by heating.
• H2O needs to be driven off in order to convert a hydroxide to an oxide.

• The other product is MgO.
• The chemical equation is balanced.

• Some of the substances are identified in the reaction summary diagram.
• It's easier to identify other reactants and products when you write down an incomplete chemical equation.

• Use Figure 21-19 to write chemical equations for the sequence of reactions in which borax is converted to diborane.

• Use Figure 21-19 to write chemical equations for the sequence of reactions in which borax is converted to BF3.

• In their appearance and physical properties, aluminum, gallium, indium, and thallium are metallic.
• Table 21.6 contains the properties of the group 13 metals.

• aluminum is the most important of the group 13 metals.
• Most of the other main-group metals are active.
• Because it is easy to oxidize to the +3 ion, aluminum is an excellent reducing agent.

• Some drain cleaners are a mixture of NaOH and Al.

• They oxidize powdered aluminum in highly exother water when they are added to Air or other oxidants.

• The evolved H21g2 helps 2 yield Al2O3: plug a stopped-up drain.

• Al2O31s2 C/rH is a good reducing agent because it will extract oxygen from grease.

• The thermite reaction is visually stunning.
• There is a photograph in the margin.

• Gallium metal is used in electronics.
• It is used to make GaAs, a compound that can convert light into electricity.
• This semiconducting material is also used in solid state devices such as transistors.

• Indium is a silvery metal.
• InAs can be used in low-temperature transistors and as a photoconductor in Richard Megna/Fundamental Photographs optical devices.

• Thallium and its compounds have few industrial uses.
• It is possible to use high-temperature super conductors.
• A thallium-based ceramic with the approximate formula Tl2Ba2Ca2Cu3O8+x exhibits superconductivity at temperatures as high as 125 K.

• The +3 oxidation state in the compounds of aluminum makes it a superconducting material.
• Gallium favors the loses electrical resistance and oxidation state.
• Under a certain temperature, indium compounds can be found.

• This preference is usually reversed in thallium.
• The oxide Tl2O, the hydroxide TlOH, and the carbonate Tl2CO3 are formed by thallium.
• The compounds are ionic and in some respects above 0 K.

• TlOH is a strong base in the solution.
• The Tl+ 3Xe44f145d106s26p1 is formed by Thallium.

• The inert pair is the pair of electrons.
• The post-transition elements have the electron configuration 1n - 12s21n - 12p61n - 12d10ns2.
• The small bond energies and lattice energies associated with the large atoms and ion at the bottom of a group are not enough to offset the ionization energies of the ns2 electrons.

• The third most abundant element is aluminum, which makes up 8.3% of Earth's solid crust.
• The United States produces more than 5 million metric tons of aluminum each year.

• When an aluminum cap was placed on the Washington Monument in 1884, it was a semiprecious metal.
• It cost $1 per ounce to produce, equivalent to the daily wage of a skilled laborer.

• All this changed in two years.

• When a student of Le Chatelier, and Al2O31s2 is added to NaOH, it causes the solid to be dissolved.

• When the solution is slightly acidified, there is a change in color.
• Pure Al2O3 can be obtained by heating the Al1OH23.

• The high melting point of Al2O3 makes it a poor electrical conductor.
• It's not feasible without a better conducting solvent.
• Hall and Heroult discovered that.

• It is a good conductor in the molten state.
• In the Hall-Heroult process, molten cryolite is used to produce aluminum metal.

• The amount of energy used to produce aluminum is very high.
• This is more than three times the amount of energy used in the electrolysis of Na.
• Because of the high energy requirements for producing Al(s), aluminum production facilities are usually located near low-cost hydroelectric power sources.
• 45% of the Al produced in the United States is obtained from the recycling of scrap aluminum, which is less than 5% of the energy required to recycle Al.

• Dry ice was added to the Fe1OH23s2 and the 3Al1OH243-1aq2 to make them slightly acidic.

• The steel tank has a carbon lining.
• Carbon is used to make the anodes.

• In the production of Al, the electrolysis bath needs to be kept at 1000 degC, which is done by means of electric heating.
• There are two other factors involved in the large energy consumption.
• The mass of Al is 27 g mol-1.
• 9 g Al is the electric current equivalent to the passage of one mole of electrons.
• One mole of electrons can produce 12 g Mg, 20 g Ca, or108 g Ag.
• Al is an outstanding energy producer when it is used in a battery because of the same factors that make Al a significant energy consumer.

• The ionic character of AlF3 is considerable.

• The molecule is composed of two AlX3 units.

• The two metal atoms are bridged with two Cl atoms.
• If the Al atoms are sp3 hybridized, then bonding in this molecule can be described.
• There are two ways in which the Cl atom bonds to the Al atoms.
• The bond to one Al atom is a conventional covalent bond because each atom contributes one electron to the bond.

• The aluminum halides are also called Lewis acids.

• They accept a pair of electrons and form adducts.

• Adding an alkyl group to a benzene ring is the most common reaction of this type.

• There are units to form Al2Cl6.

• The cation attacks the benzene ring, freeing a proton that reacts with 3AlCl.

• The cryolite, Na3AlF6, is an important halide of aluminum.
• Natural deposits of cryolite can be found almost nowhere else.

• The O2 is an ion that is in the middle of holes.
• Ruby and Fe2+ and Ti4+ are used as abrasives in Alumina, which is a very hard material.
• It is resistant to heat and is used in linings for blue sapphires.

• Artificial gemstones are used.
• The aluminum oxide is very high in reactivity and made by fusing corundum peratures.

• A thin, impervious coating of Al2O3 protects aluminum against reaction with water in the range of 4.5-8.6.
• An aluminum object is used for the anode in a bath of H2SO41aq2.

• The Al2O31s2 + 6 H+1aq2 + 6 e Al2O3 coating can be obtained.
• The oxide can be made to absorb certain substances.
• Anodized aluminum is used to make everyday items, such as the drinking cups shown in the photograph in the margin, and is also used in architectural components of buildings, such as bronze or black window frames.

• It reacts with bases to form a cup.

• The most use of aluminum sulfate in important commercial aluminum compound is in Groups 1, 2, 13, and 14.
• The acidic hot concentrated H2SO41aq2 is prepared by the reaction of sizing paper.

• Half of the calcium carbonate used in water purification is produced in the United States.
• When aluminum sulfate is added, the water's pH is adjusted so that it maintains an alkaline medium.

• It removes suspended particles from the water.
• The size of paper is an impor tant use.
• The sizing agent is deposited in the paper.

• In the industrial world, alums are a large class of double salts.

The double salts are Li+ M1I2 and K+, Na+, or NH4

• Baking powders and potassium aluminum sulfate are used in dyeing.
• The fabric is heated in steam after being dipped into a solution of alum.

• In both cases, the cations are sufficiently polarizing.

• The hydroxides of aluminum and beryllium can be found in basic solutions.

• Both metals form a strong oxide coating in the air.

• The C4 Be2C and Al4C3 ion are contained in the metals.
• Lewis acids and Friedel-Crafts catalysts can be created by Be and Al form halides.

• If KF is present, CONCEPT ASSESSMENT AlF is almost insoluble.

• To make something happen.

• Tin and lead have metallic properties.

• Silicon is classified as a metalloid but is mostly nonmetallic in its chemical behavior.
• Semiconductor behavior is also exhibited by Silicon.
• Carbon is a nonmetal.
• We will talk about carbon, Silicon, tin and lead, but germanium is not mentioned.

• The most striking differences between carbon and Silicon are outlined in Table 21.7 in the periodic table.
• Carbon-atom chains and rings play a central role in establishing the chemical behavior of carbon.

• The focus of organic chemistry and biochemistry is the study of the chains and rings.

• A Catenation is the joining of atoms into chains.

• The chemistry of carbon is emphasized in this section.

• Some of the carbon that is distributed in Earth's crust is rich enough for commercial exploitation.
• The majority of industrial graphite is made from carbon-based materials.
• The high-carbon content material needs to be heated to a temperature of 3000 degC in an electric furnace.

• Even when dry, it has excellent lubricating properties.
• The planes of carbon atoms are held together by weak forces and can easily slip past one another.
• This property is useful in pencil lead, which is a thin rod made from a mixture of graphite and clay that glides easily on paper.
• It is used for its ability to conduct electric current, and not the other way around.

• The ability to tolerate high temperatures is what determines the use of Graphite in high-temperature environments.

• The Bruce H. Frisch/Science Source weight composites are made with a mixture of graphite fibers and fabric.
• These materials are used in a wide range of products.
• When carbon-based fibers are fibers.

• The more stable form of carbon is diamond.

• The metal is usually mixed with the substance.
• The liquid metal is converted to diamond when the metal is melted.
• Diamonds can be picked out of the metal.

• At room temperature and pressure, we might expect diamond to return to its original form.
• Fortunately for the jewelry industry and for those who treasure diamonds as gems, many phase changes that require a rearrangement in bond type and crystal structure occur extremely slowly.

• That is the case with the diamond-graphite transition.

• The point is marked by an arrow.

• Natural diamonds are used as gemstones.
• Synthetic diamonds are shown in the margin for industrial purposes.
• The two key properties are used in the industrial use.
• Diamonds are very hard and are used as abrasives.
• There is no harder substance.
• Diamonds have a high thermal conductivity, so they are used in drill bits for cutting steel and other hard materials.
• The lifetime of the bit is increased by the rapid dissipation of heat.
• Because of their expected properties of diamond to the metal, diamond films can be deposited directly onto metals.
• When a metal is coated with a diamond film, the resulting material has a high thermal conductiv films.
• The journal has used such materials in heat sinks for computer chips.

• The fullerenes earned that honor in 1991.

• Mixed crys talline or amorphous structures are some of the forms of carbon that can be obtained.
• A smoky flame can be caused by incomplete combustion of natural gas in a Bunsen burner.

• Carbon black is used as a material in rubber tires, as a material in printing ink, and as a transfer material in carbon paper, typewriter ribbons, laser printers, and photocopying machines.
• New allotropes of carbon have been isolated.
• The allotropes were presented in Chapter 12.

• The molecule C60 has a shape similar to a soccer ball and is remarkably stable.
• C70, C74, and C82 are other fullerenes.
• The production of enes can be done by laser under a helium atmosphere.
• soot does not contain fullerenes because nitrogen and oxygen interfere with the process of forming them.

• We talked about various forms of carbon in Chapter 12.
• In this chapter, we focus on Graphene.

• The synthetic diamonds are called graphene.
• A sheet of Graphene is rolled into a cylinder.
• A spherical ball is formed when the flat sheet to pucker is replaced by pentagonal rings.

• Graphene has interesting electronic properties because the electrons in the sheets are moving very quickly.

• Graphene is expected to play a role in the development of electronic devices.

• It was thought impossible that a sheet of carbon could be made.
• In 2004, scientists in the United Kingdom used a technique called micromechanical cleavage to isolated graphene.

• The process is similar to drawing with a pencil, with the "lead" of which isgraphite, and looking at the traces left by the pencil.
• Another way to get Graphene is to peel away layers of carbon atoms from a Graphene surface using a process called exfoliation.
• The methods used to produce the flakes contain up to 10 layers of Graphene.
• The single-layer flakes have to be found among the thicker ones.

• Coke and charcoal are carbon-based materials.

• In blast furnaces, it is used to reduce iron oxide to iron metal.

• CO and CO2 are the chief oxides of carbon dioxide.
• The amount of CO2 in the air is around 400 parts per million.

• It occurs to a lesser extent.
• The two oxides are important in many ways.

• A fuel-lean mixture is burned in an automobile engine.

• A fuel-rich mixture is burned in an automobile engine.

• Fossil fuels in automobile engines cause CO to be an air pollutant.

• Carbon monoxide binding to the iron atoms in hemoglobin is stronger than oxygen.
• Carbon monoxide's toxicity arises because it prevents hemoglobin from binding with oxygen.
• The molecule shown here is called a heme group.
• Four nitrogen atoms surround an iron atom in the center of the group.
• In hemoglobin, an O2 molecule projects above the plane of the iron and nitrogen atoms, but here it has been replaced by a CO molecule.

• Air pollution can be caused by incomplete combustion of gasoline and a loss of efficiency.
• If CO1g2 is formed as a combustion product, gasoline will evolve less heat.

• Carbon dioxide can be obtained directly from the atmosphere, but it is not an important source.
• The sources of CO2 are summarized in Table 21.8.

• Dry ice is the main form of carbon dioxide used for freezing, preserving, and transporting food.
• Carbonated beverages make up 20% of CO2 consumption.
• Oil recovery in oil fields is an important use.
• The major use is by plants and algae.

• Green plants use atmospheric CO2 as a source of carbon-containing compounds.
• There are major exchanges between the surface of Earth and the atmosphere.

• Determine how much less heat is produced per mol of C burned in reaction than in reaction.

• The details of the process have been known for a few decades.
• His research on C6H12O6 involved up to 100 sequential steps for the conversion of 6 mol CO2 to 1 mol in 1961.
• The assimilation of carbon dioxide in plants is a representation of the overall change.

• The overall reaction is very cold.
• The required energy comes from the sun.
• Plants have a green color called chlorophyll.

• The reaction produces atmospheric oxygen.

• Animals pass carbon atoms to plants.
• When the animals breathe and expel gas, some carbon is returned to the atmosphere as CO2.
• As plants and animals die and their remains are broken down, additional CO2 returns to the atmosphere.
• Coal, petroleum, and natural gas are converted to carbon in decaying organic matter.
• This carbon can't be used for photosynthesis.

• The cycle of CO2 through the oceans is not represented in the drawing.
• The small floating green organisms convert CO2 to organic compounds.
• All the animals in the oceans are supported by Phytoplankton, which are at the bottom of the ocean food chain.

• Huge quantities of carbon have been deposited in carbonate rocks.
• These come from the shells of dead mollusks.

• Human activities are more important in the carbon cycle than they were in the past.
• Fossil fuels are burning more carbon dioxide than stored carbon.
• A future global warming and an increased level of atmospheric CO2 are possible consequences of this distortion of the carbon cycle.

• The natural carbon cycle has become a topic of debate.

• The synthesis of NH3 is dependent on the use of the reforming of natural gas.

• There are three main uses of carbon monoxide.
• One is making other compounds.

• CO can be used as a reducing agent.

• The reaction can be done by heating coke and Fe2O3 in a blast furnace.

• A third use of CO is as a fuel, usually mixed with other com bustible gases.
• Section 7 discussed this.

• The carbides are ionic.

• Carbon disulfide is a highly volatile liquid that acts as a solvent.
• It's uses as a solvent are decreasing because it's poisonous.
• The manufacture of rayon and cellophane is an important use.

• CCl4 has been extensively used as a solvent, dry-cleaning agent, and fire extinguisher, but these uses have been declining because CCl4 is a known carcinogen.

• Some groupings of atoms have characteristics of a halogen atom.
• HCN is a liquid that can boil at room temperature.
• It is a very weak acid.

• HCN has uses in the manufacture of plastic.

• It is used in organic synthesis, as a fumigant, and as a rocket propellant.

• Silicon is the second most abundant element in the Earth's crust.
• Silicon is to the living world as carbon is to the mineral world.

• When coke is used in an electric arcs furnace to reduce the amount of 1SiO22 in the sand, Silicon Elemental Silicon is produced.

• Si + 2 CO1g2 is a very high purity Si for solar cells.
• The by-product 2 is the Na2SiF6 required for this process.

• The Si atom is surrounded by four O atoms.

• The way materials scientists represent silicates and similar materials is the same as this view.

• In the manufacture of transistors and other Semiconductor Devices, high-purity Silicon is required.

• The only stable oxide of Silicon is SiO2.

• It is a network covalent solid.
• Figure 21-31(a) shows the structure of a network covalent solid.

• The structure is similar to the diamond structure and has some similarities to diamond.
• The raw material for the glass and ceramics industries is sibel.

• There are a number of ways in which these tetrahedra can be arranged.

• The empirical formula is LiAl1SiO322.

• There are two corners to only two.
• The mineral has a fibrous appearance.
• The formula of the mula is empirical.

• Each Si atom is joined into three adjacent Si atoms.

• Bonding between the sheets is not as strong as it is between the sheets.

• This is the most common arrangement in the majority of silicate minerals.

• SiO2 is a weakly acidic oxide.

• The H atom and one OH group Silicate anions are bases and can be acidified.
• Crystalline solid or powder that is not of H2O are eliminated from the sample.

• These hydrated silicates are formed by the elimination of water from the neighboring molecule of silicic acid.

• We might expect carbon to form oxides with similar properties because they are both in group 14 of the periodic table.

• The second- and third-period members of group 2 were contrasted on page 994.
• The carbon member of the group is different from the higher period members.

• The bridges are made of Si.

• There are four rings of nonoxygen atoms in the b-cage.

• The arrangements described above have an important consequence.

• zeolites have important properties.
• When eight b-cages are joined together by sharing example, zeolites have been used as sieves to remove rings.
• The structure of the mixture is called a structural unit.
• In the past, zeolites have been used to remove water from gases.
• For occurring.
• The four-membered rings of the zeolite can be separated from the benzene and regenerated by heating.

• As an ion exchange material, zeolites is an important application.
• The structure of the water is more open than the Ca2+, Mg2+, or Fe2+ concentrations.
• The structures can be used to exchange the ion with the Na+ ion.

• The cations from water react with CO3 or anions of soaps to form insoluble precipitates.

• The formation of boiler scale low zeolite can be obtained by heating the water and building up the zeolite in the pipes or under the vacuum.
• The containers are used for boiling water.

• When there is high affinity for water, the equation below represents the exchange.

• A cation-exchange resin is shown here.

• By the time the water reaches the bottom of the column, all the multivalent ion have been removed and only Na+ ion remain as counterions.
• The exchange can be represented as 2 NaR + M2+ D MR2 + 2 Na+.
• The reaction takes place in the forward direction.
• The reverse reaction is favored when the concentration of NaCl(aq) is high.

• The formation of insoluble precipitates is prevented by the replacement of Ca2+ ion by Na+ ion.

• The zeolites are used in detergents to help remove Ca2+ and Mg2+ ion that may be present in water used for washing clothes.
• The removal of these ion helps detergents foam better and prevents the formation of insoluble calcium and magnesium compounds.

• Some automobile engines need fuels with low boiling points and some need fuels with branched-chain hydrocarbons.
• Short-chain or branched-chain hydrocarbons can be converted using zeolite catalysts.
• The role of zeolites in the industry is important.

• Glass hydrated silicate polymers are important in the ceramics industry.

• The final ceramic product is processed into the gel.
• The sol-gel process can produce lightweight ceramic materials.

• Applications that take advantage of the ceramic's mechanical and structural properties at high temperatures are included in the general category.
• These properties have been explored in the development of ceramic components for gas turbine and automotive engines.

• A liquid mixture of sodium and calcium silicates can be created if salt and calcium carbonates are mixed with sand.

• The structural units in glass are not in a regular arrangement.
• The melting behavior of a glass and a solid is different.
• A glass will melt over a wide temperature range, whereas a solid will melt at a specific point.
• The methods of making different types of glass are described later in this section.

• The direct reaction of Si and CH3Cl is typical.

• Silicones are important because they are versatile.

• Silicones can be obtained as oils or rubber-like materials.
• Silicone oils do not oxidize when heated.
• They can be cooled to low temperatures without becoming silicones.

• Silicone oils are good at high temperatures.
• hydrocarbon oils break down at high temperatures and then solidify at low temperatures.
• Silicone rubbers retain their elasticity at low temperatures.
• They are useful in caulking around windows.

• Silicon reacts with the 1X22 to form SiX4.
• At room temperature, SiF4 is a gas, SiCl4 is a liquid, and SiI4 is a solid.
• Both SiF4 and SiCl4 can be hydrolyzed with water.

• It is possible to make very pure Silicon for transistors used in computer chips, as well as very finely divided Silicon for use as a reinforcing filler in Silicone rubber.

• The art of glassmaking has been around for a long time.
• There are beautiful stained-glass windows in medieval and modern churches, and ancient glass containers for perfume and oil in many museums.
• Glass is used in almost every facet of life.

• The mixture can be fused at a relatively low temperature compared to the melting point of pure silica, and it is easy to form into shapes.
• The glass can be used for things like drinking glasses and windows because of the effect of the calcium and sodium ion.
• The ultimate glass product is a mixture of sodium and calcium silicates.

• There are stained-glass windows in the Chapel of Thanksgiving in Dallas, Texas.

• The Chemistry of the Main- Group Elements I: Groups 1, 2, 13, and 14 Glass has a distinctive green color.
• The process of making glass can be made simpler with the addition of MnO2.
• The violet color is caused by the oxidation of green FeSiO3 to yellow Fe21SiO323 and the reduction of Mn2O3.
• The yellow and violet are the same color.
• CoO can be used to impart color where desired.
• Additives such as calcium are used to make an opaque glass.
• A glass with exceptional transparency can be made by incorporating lead oxides.

• The dimensions of soda-lime glass change with temperature.
• The glass can't survive thermal shock.
• The lanterns used in the early days of railroads were limited by this limitation.
• The hot glass in the lanterns would shatter in the rain.
• The problem was solved by adding B2O3 to the glass.
• Pyr glass is used in cookware in the home and in chemical laboratories.

• The distorted images produced by the thick bottoms of drinking glasses are caused by small bubbles or impurities in most glass.
• Light cannot be transmitted over long distances without distortion or loss of signal.
• A special glass made of pure silica is required.
• A series of chemical reactions can be used to make this glass.
• First, coke is used as a reducing agent.
• SiCl41g2 is formed when the Silicon is allowed to react with Cl21g2.
• The Si4 is burned in a flame.
• SiO2 is a fine ash, and chlorocarbon compounds are gaseous products.
• The SiO2 can be melted and drawn into the fine fibers of the cable.
• The United States produces millions of kilometers of fiber-optic cable each year.

• Boron is a solid acidic oxide.
• Al2O3 is amphoteric and CO2 is acidic.

• Boric acid is a weak acid.

• There is a wide range of borates and silicates.

• This diagonal relationship is not easily understood and can't be seen in terms of charge density since the bonding in boron compounds and in Silicon compounds is exclusively covalent.
• Both metalloids have similar electronegativities and have the same size.

• The data in Table 21.9 suggests that tin and lead are similar.

• They are soft and melt at low temperatures.

• They have the same tendencies to be oxidation to the +2 oxidation state.

• Tin and lead can be found in two oxidation states, +2 and +4, which is an example of the inert pair effect.
• In the +2 oxidation state, the ns2 is not involved in bond formation, whereas in the +4 oxidation state, they are.
• Tin has a tendency to be in the + 4 oxidation state.
• The lower oxidation state is favored farther down a group in that trend.

• tin and lead both have two common forms, but lead has a single solid form.
• The b (white), or metallic, form of tin is stable above 13 degC.
• When a sample of b tin is cooled, it must be kept below 13 degrees for a long time before it becomes a tin.
• The transformation takes place quickly and with dramatic results once it begins.
• The tin expands and falls to a powder because it is less dense than the b variety.
• The objects made of tin are destroyed by this transformation.
• Some organ pipes are made of tin or tin alloy, which can cause a problem in churches in cold climates.

• Tinplate and plat ing iron are used in cans for storing food.
• The low-melting alloys are used to join wires.
• Organ pipes are made out of Sn and Pb.

• The coke is used to make the metal.

• Half of the lead produced is used in batteries.

• To protect against X-rays, other uses include the manufacture of solder and other alloys.

• Tin and lead exhibit the +2 and +4 oxidation states.

• Table 21.10 shows the charge densities of the ion.
• Many of the compounds containing tin in the +2 oxidation Lead Ions state are covalent; however, a few ionic solids containing the Sn2+ ion are known.
• Lead is found in many ionic solids.

• Tin forms two primary oxides.

• It is possible to use SnO2 as an abrasive jewelry.

• Some of the chemistry of lead oxides is not fully understood.
• There are several batteries, glass, ceramic glazes, cements, metal-protecting paints and lead oxides used in the manufacture of lead-acid.
• Other lead compounds are usually made.

• Because lead tends to be in the +2 oxidation state, lead(IV) compounds tend to undergo reduction to compounds of lead(II) and are therefore good oxidizing agents.
• PbO2 is a case in point.
• It was noted in Chapter 19 that it was used in lead-acid storage cells.

• Pb21Claq2 + 2 H2O1l2 + Cl21g2 Edegcell is a compound.
• There is a yellow oil that reacts with the air in a way similar to the way lead reacts with water.

• Tin and lead are not the same.
• Lead(II) chloride is a white insoluble ionic solid, while tin(II) chloride is a covalent solid.
• The molecule is a V-shaped one in the gas phase.
• Because of the lone pair, we might expect the base to be named after Lewis.
• It acts as a Lewis acid.
• For example, reacts with a substance to form a substance.

• There are important uses for both tin-cl2 and tin-cl4.
• Tin(II) chlo ride is used in the quantitative analysis of iron ores to reduce iron(III) to iron(II) in a solution.
• The form of tin recovered from scrap tinplate is called Tin(IV) chloride.
• Tin(II) fluoride was used as an anticavity enhancer to toothpaste but has largely been replaced by NaF in gel toothpastes.

• One of the fewsoluble lead compounds is Pb1NO322.

chrome yellow is a yellow color caused by lead(II) chromate PbCrO4 being added to Pb1NO3221aq2

• Lead has been used in plumbing systems since the ancient Romans.
• Lead can be found in cooking and eating utensils and pottery glazes.
• Lead poisoning was the cause of "dry bellyache" for some North Carolina residents who drank rum from New England.
• The equipment used to make the rum was made of lead.

• Depression and nervousness are caused by mild forms of lead poisoning.
• More severe cases can cause permanent damage.
• Lead causes the heme group in hemoglobin to be disrupted.
• The effects of Pb>dL in blood can be seen in small children.
• The drop in lead levels in blood is a result of the phaseout of leaded gasoline.
• The graph shown in Figure 21-35 shows a decline in blood lead levels as well as a decline in the use of lead in gasoline.
• Lead-based painted surfaces in old buildings and soldered joints in plumbing systems are the main sources of lead contamination.
• Lead has been eliminated from modern plumbing solder.
• The disposal of lead is closely monitored.

• Most of the lead metal production is provided by recycling.

• There was a decline in the level of lead in the blood of a representative human population, just as there was a decline in the use of lead in gasoline in the 1970s.

• The metallic tin is kept in contact with the Sn2 to prevent air oxidation.
• This contact helps prevent oxidation.

• GaAs may be one of the most versatile high-tech materials of our time.
• Chapter 21 of Gallium Arsenide has a feature called "Focus On" on the MasteringChemistry site where you can discuss some of its properties.

• The bonding in affinity, electronegativity, and diborane are described by the trends in atomic or ionic radii, ionization energy, electron two-electron bonds.

• The principal metal of group 13 is aluminum, which is the most active of the metals, as indicated by the fact that large-scale use is made possible by an effective their low ionization energies and large negative electrode method of production.
• The basis for separating Al2O3 from magnesium in group 2 is the same as the basis for separating Al2O3 from impurities.
• Most of the alkali metals are Fe2O3.
• molten Na and molten salts are used to carry out lysis.

• It's very reactive.
• A Lewis acid and an anionic sulfate group are attached to each other by a long-chain hydrocarbon termi molecule.

• Some group 2 metals are prepared by the 21-5 Group 14: The Carbon Family-- Group 14 electrolysis of a molten salt and some with one nonmetal, two metalloids and chemical reduction.
• Group 14 is notable for the carbonates, especially the differences between the first two members.

• The proper pounds of the form species include the anion 3MX ties exhibited by the pairs of elements.

• One nonmetal, B, and the metals Al, Ga, In, and Tl can be found in the Boron Family.

• The elements sieve and treat hard water.
• Tin and lead can be obtained by reducing their oxides.
• The metal has some similarities.
• The soft met lurgical method of converting a sulfide to an oxide has low melting points.

• If you don't perform detailed calculations, you can show that the reaction correctly describes the dissolving action of rain on limestone.

• The dissolution of cal is described by the relevant equations.

• Carbonic acid is 2 in water.

• The carbonate 2O D H3O+ + CO3 K ion concentration observed in a carbonic acid solution is much greater than the CaCO3 solution.

• The carbonate ion from the dissolution of CaCO3 acts as a common ion in the carbonic acid equilibrium when the two processes occur.

• The qualitative correctness of the reaction has been assessed.
• It is more difficult to calculate the quantitative extent of the dissolution of CaCO31s2 in rain.
• The partial pressure of atmospheric CO2 in equilibrium with rain affects the calculation.
• Data is given in Chapter 16, Practice example A.

• When NaCN and Al1NO323 are dissolved in water, write chemical equations for the reactions that occur.
• Data from Appendix D can be used to explain why a precipitate of Al1OH23 forms when equal volumes of NaCN and Al1NO323 are mixed.

• CaSO41s2 + 2 NaOH1aq2.

• Write a net ionic equation.

• molten NaOH was used as the electrolyte.

• A 1.26 L sample of KCl(aq) is electrolyzed for 3.50 minutes and has a current of 0.910 A.

• The power needed to regulate the heartbeat is 5.0 mW.

• The starting material is outlined in the diagram.
• The final product of the reac is Mg metal.

• The principle of conser indicated substances seems to have been violated by this process.

• As written, Mg1HCO3221s2 is heated to a high temperature.
• Appendix D is needed for Ba21l2.

• The formula B4H10 is used for the molecule tetraborane.

• Lewis structures should be written for the following species.
• The metal sample in the testing can be destroyed.

• Write chemical equations to represent an event.

• The oxidation states of the atoms in a perborate obtained aluminum can be assigned based on the structure of page 1003.

• To represent the reactions, write chemical equations.

• The compounds do not exist because of the ex reaction.

• In some foam-type fire extinguishers, the reactants give very unstable B8F12 and BF3.
• The B8F12 molecule is found in Al21SO4231aq2 and NaHCO31aq2.
• When the extin is unstable, the reactants mix and produce a ring of boron atoms.
• Write a balanced chemical equation.

• To rep B8F12, write a net ionic equation.
• Two of the six units in the B8F12 resent this reaction.

• Gallium trichloride is a very active catalyst.
• When water is added to this mix.
• The compounds CO21g2 and Al1OH231s2 are two solid state compounds and GaCl3 is a dimer with the formula of the products.
• Draw a structure for the two products.

• To represent from PbO2, write a chemical equation.

• PbO2 is a good oxidizer.

• Carbon tetrachloride and S2Cl2 can be formed by 10-4 M2 to MnO4.

• Aqueous tin(II) ion is a good reducing agent.

• There is further reaction of carbon disulfide and S2Cl2 pro use data.
• Write Sn2+(aq) is a good reducing agent that can be used to reduce a series of equations.

Would you expect the silanol molecule to react to H2O?

• A chemical that should exist as a solid is to melt and maintain molten NaOH.
• Commercial process for the production of on a storeroom shelf is preferred for the mixture of solid and liquid in a container.
• Give a plausible reason for that.

• Give a reason for the discrepancy.

• The ion is stable in only a small piece of dry ice when added to a solution.
• CsI3 is stable with respect to Ca1OH221aq2.

• To explain the ion, write chemical equations and suggest a reason why CsI3 is stable.

• The melting point of NaOH is 322 degC.

• The electrolysis of 0.250 L of 0.220 M MgCl2 is a metal ion.
• A mixture of H2 and water energy changes and the enthalpy changes for the vapor is collected at 23 degrees.
• If the process is carried to C/ hydrGdeg and C/hydrHdeg respecively, will it be marked by Mg1OH221s2?

• LiO21s2 has never been isolated.

• An aluminum production cell of the type pictured in LiO21s2 and assess whether LiO21s2 is thermodynam Figure 21-24 operates at a current of 1.00 * 105 A and ically stable with respect to Li2O1s2 and O21g2.

• The Kapustinskii equation can be used to produce chemical change.

• The mass of Al can be produced by this cell.

• If the electrical energy is needed to power this cell.

• The process O21g2 + e- is used in a power plant that has an effi of 43 kJ mol-1.

• If the dation of the anode to CO21g2 permits LiO21s2 to be stable with respect to the electrolysis to occur at a lower voltage than if Li2O1s2 and O21g2, then you should use your result from part (c) to 2 Al1l2 should be neglected.

• A saturated solution of Pb1NO322 product is obtained.
• All the Ca appears in has a relative humidity of 97%.

• A balanced equation is needed for this reaction.

• Each case can have the dissolution of MgCO31s2 in NH.

• The crystal structure was deduced by Li.

• The relationship of elec water contains trace amounts of magne trode potentials.

• Expressed as grams of salt per liter.

• A set of three steps for the reduction of water should be estimated.
• How does your estimate compare to the H+?

• Lake water calls for a pinch of borax.

• They are at the site of underwater solutions.

• If not for the fresh water entering through the Sierra Nevada mountains, the lake level would be lowered by three meters per year.
• The salts in the lake are the chlorides, bicarbonates, and sulfates of sodium.

• Na is being used as the reducing agent.

• A chemist knows that CaSO4 2 H2O, with ammonium carbonate to toward oxygen than is iron, is more reactive than iron.
• There was evidence provided in the water.

• Write chemical equations to represent the strongest reducing agent in each outcome.
• Why is aluminum likely to happen if it is so reactive?

• There are several substances listed.
• The following descriptions are given in a chemical dictionary.
• The pair reacts with water.
• The chemical equations are based on the descriptions.

• Balance and complete the following.
• So state if there is no reaction.