18.4 Structure and General Properties of the Nonmetals
The periodic table has elements located in the upper right portion.
Their properties and behavior are very different from those on the left side.
More than half of the nonmetals are gases, one is a liquid, and the rest are soft and hard.
The nonmetals have a lot of chemical behaviors.
They form many different ionic and covalent compounds by including the most reactive and least reactive elements.
An overview of the properties and chemical behaviors of the nonmetals is presented in this section.
Nonmetals are important in biological systems.
Trends in electronegativity allow us to predict the type of bonding and physical states in compounds involving nonmetals.
We know that electronegativity decreases as we move down a group and increases as we move up a group.
The nonmetals have higher electronegativities than do metals, and compounds formed between metals and nonmetals are generally ionic in nature.
The metals form cations and the nonmetals form anions.
On the other hand, compounds formed between two or more nonmetals have small differences in electronegativity between the atoms.
These substances are gases, liquids, or volatile substances and can be found at room temperature and pressure.
Nonmetals do not form monatomic positive cations in normal chemical processes because of their high ionization energies.
The selenide ion is one of the monatomic nonmetal ion that are anions.
Remember that an element exhibits a positive oxidation state when combined with a more negative element and that it exhibits a negative oxidation state when combined with a less negative element is a common oxidation state.
These oxidation states are found in ionic and covalent compounds.
The first member of a nonmetal group has different behaviors than the other group members.
The smaller size of the first member of each group allows better overlap of atomic orbitals.
P bonds are formed between nonmetals that are the first member of a group.
Stable p bonds to itself are not normally formed by sulfur.
The variety of oxidation states displayed by most of the nonmetals means that many of their chemical reactions.
Most metals oxidize.
The oxidation state of the metal becomes positive as it undergoes oxidation and the nonmetal becomes negative as it undergoes reduction.
The strongest oxidizers within their groups are ferriine and oxygen.
Group 17 is the strongest oxidizing agent.
A nonmetal oxidizes an element that lies to its left.
The more difficult it is to oxidize the anion formed by the nonmetal, the stronger it is.
The most stable negative ion are formed by elements at the top of the group.
The strongest oxidizers are ferriine and oxygen.
Oxygen only exhibits a positive oxidation state when combined with fluorine.
All nonmetals form compounds with oxygen, yielding covalent oxides.
The oxides react with water to form oxyacids.
An oxyacid is an acid consisting of hydrogen, oxygen, and some other element.
Notable exceptions are carbon monoxide, CO, nitrous oxide, N2O, and NO.
The product is an acid.
The oxides that do not exhibit one of their common oxidation states react with water.
The nonmetal is reduced in these reactions.
The discussion on acid-base chemistry can be found in the chapter.
The hydrogen compounds of the nonmetals exhibit an acidic behavior in the water.
The acid strength of the nonmetal hydrogen compounds varies from left to right.
Chapter 18 contains Representative metals, metalloids, and nonmetals.
H2O is a weaker acid than H2S and H2Se.
The structures of metals are very similar to those of nonmetals.
The metals in the tightly packed array do not have any bonds.
Many nonmetals are composed of individual molecules.
There is delocalization of the electrons throughout the solid in nonmetals.
The noble gases are all monatomic, whereas the other nonmetal gases are diatomic.
The other halogens are diatomic as well, with I2 being a solid under normal conditions.
The increasing strength of intermolecular London forces with increasing molecular mass and increasing polarizability can be seen in the changes in state as one moves down the halogen family.
Oxygen has three allotropes: O2, dioxygen, and ozone.
White, red, and black are the colors of Phosphorus.
Sulfur has many allotropes.
There are many carbon allotropes.
Less people know about the recent discovery of fullerenes, carbon nanotubes, and Graphene than they do about diamond, graphite, and charcoal.
The physical properties of three nonmetals are described.
Carbon can be found in many forms, such as diamond, graphite, charcoal, coke, carbon black, and fullerene.
The diamond is a giant molecule.
The crystals are very hard and have high melting points, because carbon-carbon single bonds extend throughout the crystal to form a three-dimensional network.
The structure consists of layers of carbon atoms with each atom surrounded by three other carbon atoms.
There are many resonance forms that can be found in the OpenStax book at http://cnx.org/content/col11760/1.9.
The s and p bonds are strong, but the forces between layers are weak.
The layers are held together by London dispersion forces.
See the discussion of the weak forces in the chapter on liquids.
The soft, flaky character that makes it useful as the so-called "lead" in pencils and the slippery character that makes it useful as a lubricant are the result of the weak forces between the layers.
The p bonds hold electrons that can move throughout the solid and are responsible for the electrical conductivity of the material.
Carbon black, charcoal, and coke are other forms of carbon.
Carbon black is a form of carbon created by incomplete burning of natural gas.
In the absence of air, it is possible to make coke and charcoal by heating wood and coal.
There are new forms of carbon in the soot generated by a smoky flame and in the vapor produced by the high temperatures in a vacuum.
One of the new forms was isolated by Professor Richard Smalley and coworkers at Rice University.
The structure of C60 is icosahedral.
Two allotropes of carbon are graphene and carbon nanotubes.
There is a relationship between the forms.
Representative metals, metalloids, and nonmetals Graphene is a very strong, lightweight, and efficient conductor of heat and electricity.
The system is stable because of resonance.
There is no stacking of the layers to give a three-dimensional structure.
The University of Manchester won the 2010 Nobel Prize in physics for their work on Graphene.
The simplest way to prepare Graphene is to use a piece of tape to remove a single layer of Graphene from a piece of Graphene.
The method works because there are weak dispersion forces between the layers.
Alternatively, you can deposit a single layer of carbon atoms on the surface of some other material, or you can make it at the surface of Silicon.
There are no commercial applications of Graphene.
Its unusual properties, such as high electron mobility and thermal conductivity, should make it suitable for the manufacture of many advanced electronic devices.
The carbon allotropes have a cylindrical structure.
The layers wrap into a tube and bond together to produce a stable structure.
The walls of the tube are made of atoms.
Diamond is harder than carbon nanotubes.
It can be a conductor or a Semiconductor.
The conducting form is better for some applications.
The creation of carbon atoms in a vacuum is the basis for the synthesis of carbon nanotubes.
It is possible to produce carbon atoms by using a variety of methods.
The strength of carbon nanotubes will eventually lead to some of their most exciting applications, as a thread produced from several nanotubes will support enormous weight.
The current applications only use bulk nanotubes.
The mechanical, thermal, and electrical properties of the bulk material can be improved with the addition of nanotubes.
Some bicycle parts, skis, baseball bats, fishing rods, and surfboards are made with nanotubes.
Scientists noted that when phosphorus was first isolated, it glowed in the dark and burned when exposed to air.
There is only one member of the group that does not occur in the uncombined state in nature, and that is Phosphorus.
We will look at the two forms: white and red.
White phosphorus is a white, waxy solid that has a boiling point of 280 degrees.
As a solid, as a liquid, as a gas, and in solution, white phosphorus is P4 molecule with four phosphorus atoms at the corners of a regular tetrahedron, as illustrated in Each phosphorus atom bonds to the other three atoms in the molecule by single covalent bonds.
The most toxic allotrope is white phosphorus.
In the absence of air, heating white phosphorus to 270-300 degC yields red phosphorus.
Its structure appears to contain threedimensional networks of P4 tetrahedra joined by P-P single bonds.
White phosphorus is dissolved by red phosphorus.
P4 is formed when red phosphorus is heated.
The allotropy of sulfur is more complex than any other element.
In the Bible and other places, sulfur is referred to as a brimstone, and in recent times it has been discovered that it is a component of the atmospheres of Venus and Io.
The rhombic sulfur, also known as yellow, is the most stable and common allotrope of sulfur.
All other allotropes reverting at room temperature are made of Rhombic sulfur.
There are long needles of monoclinic sulfur when cooling this liquid.
The form is stable from 96 degC to the melting point.
It gradually becomes rhombic at room temperature.
The sulfur atoms are bonds to their neighbors by single bonds.
The four sulfur allotropes show eight rings.
Here are individual S8 rings, S8 chains formed when the rings open, S8 chains formed by adding sulfur atoms to S8 chains, and part of the very long sulfur chains formed at higher temperatures.
When rhombic sulfur is melted, the straw-colored liquid is quite mobile because of the spherical nature of the S8 molecule.
S-S bonds in the rings break as the temperature rises.
These chains form long chains that tangle with one another.
It does not pour easily because of the dark color of the liquid.
The dark red color is due to the electronic structure of the sulfur atoms that are dangling at the ends of their chains.
The light is absorbed differently and results in a different color.
Plastic sulfur is a rubberlike mass produced by cooling the liquid.
The formula S2 is a paramagnetic molecule like O2 with a similar electronic structure and a weak sulfur-sul.
An important feature of the structural behavior of the nonmetals is that the elements usually occur with eight electrons in their valence shells.
If necessary, the elements form bonds to make up for the missing electrons.
Members of group 15 have five valence electrons and only need three more to fill their shells.
The elements form three bonds in their free state, with triple bonds in the N2 molecule or single bonds to three different atoms in arsenic and phosphorus.
Oxygen forms a double bond in the O2 molecule, and sulfur, selenium, and tellurium form two single bonds in various rings and chains.
Each atom has a single bond in the diatomic molecule formed by the halogens.
The electron is required to complete the octet.
The noble gases don't form bonds with other noble gas atoms because they have a filled outer shell.
The most abundant element in the universe is hydrogen.
The sun is composed of hydrogen.
Most of the atoms in the universe are hydrogen atoms.
There are more compounds with hydrogen than with any other element.
The most abundant compound of hydrogen is water.
Many minerals, sugars, fats, oils, alcohols, acids, and thousands of other substances are made from hydrogen.
H2 is a odorless, tasteless, and nonpoisonous gas that can be found at ordinary temperatures.
Unlike other elements, hydrogen is composed of protium, 1H, deuterium, 2H, and tritium 3H.
There is one atom of deuterium for every 7000 H atoms and one atom of radioactive tritium for every 1018 H atoms in a naturally occurring sample of hydrogen.
The chemical and physical properties of the different isotopes are very similar, but they differ in some areas.
The vapor pressure of ordinary hydrogen is higher than that of deuterium and tritium.
The heavier isotopes are concentrated in the last part of the liquid hydrogen to evaporate.
By breaking chemical bonds, hydrogen must be prepared.
The methods of preparing hydrogen are common.
Water is the most abundant source of hydrogen.
Water gas can be used as an industrial fuel.
It is possible to produce more hydrogen by mixing the water gas with steam in the presence of a catalyst.
The water gas shift reaction is a reaction.
A mixture of hydrogen and carbon monoxide can be prepared by steam over a nickel-based catalyst.
There is twice the amount of hydrogen produced at the anode as there is at the cathode because of the diatomic nature of the elements.
The most convenient way to make hydrogen is in a laboratory.
Hydrogen gas and metal salts can be produced by metals with lower reduction potentials.
The iron and acid reaction produces hydrogen.
They are important in the inflation of life jackets.
When heated, hydrogen enters into many chemical reactions.
Two thirds of the world's hydrogen production is devoted to the manufacture of ammonia, which is used in the manufacture of nitric acid.
Hydrogen can be used as a nonpolluting fuel.
The exothermic reaction of hydrogen with oxygen releases 286 kJ of energy per mole of water formed.
The oxygen-hydrogen torch can achieve temperatures up to 2800 degrees.
The torch has a hot flame that is useful in cutting metals.
Liquid hydrogen and liquid oxygen were used in the space shuttle's engines before the fleet's retirement in 2011.
The large tank held the liquids until the shuttle launched.
hydrogen can occupy two locations in the periodic table and the 1 valence shell has a capacity for two electrons.
H+ is a group 1 element because hydrogen can lose an electron to form it.
It is possible to consider hydrogen to be a group 17 element because it only needs one electron to form a hydride ion, H-, or it can share an electron to form a single, covalent bond.
Hydrogen is a unique element that should be in the periodic table.
When heated, hydrogen reacts with the metals of group 1 and group 2.
The hydride anion, H-, is a strong reducing agent that reacts with water and other acids to form hydrogen gas.
As the electronegativity of the nonmetal increases, the reactions become exothermic and vigorous.
When heated, hydrogen reacts with nitrogen and sulfur, but it also reacts with fluorine and chlorine.
There is a mixture of hydrogen and oxygen.
To avoid the formation of an explosion in a confined space, it is necessary to exercise caution when handling hydrogen or any other gas.
Ammonia and phosphine are very weak acids and can be used as bases.
Table 18.1 contains a summary of the reactions of hydrogen with the elements.
The formation of the metal and water vapor is reduced by hydrogen.
CuO forms copper and water by passing hydrogen over it.
We will only talk about a few hydrogen compounds of the nonmetals.
Ammonia is prepared in a laboratory by the reaction of an Ammonia salt with a strong base.
Ammonia is formed when ionic nitrides react with water.
The structure of ammonia has a central nitrogen atom and three hydrogen atoms.
Ammonia is a gas with a strong odor.
This powerful odor is utilized by smelling salts.
The liquid that liquefies is a gaseous ammonia.
The enthalpy of ammonia's Vaporization is higher than that of any other liquid except water, so ammonia is useful as a refrigerant.
Ammonia can be found in water in 1 L H2O.
Ammonia is discussed in the chapter on acid-base chemistry.
The compounds of the two ion are very similar in structure and solubilities.
Ammonia is a weaker acid than water.
Ammonia is so weak that high temperatures are necessary.
The Hydrogen and This OpenStax book can be found for free at http://cnx.org/content/col11760/1.9.
The nitrogen atom in ammonia is not susceptible to reduction.
Ammonia gives NO and water.
Ammonia hot and the ion are reducing agents.
The oxidations of ammonium ion by NO - 2 and nitrate to yield pure nitrogen and N2O are of particular interest.
The replacement of one or more hydrogen atoms with some other atom or group of atoms can be used to consider derivatives of ammonia.
Inorganics include chloramine, NH2Cl, and hydrazine, N2H4.
hydrazine is a Lewis base and a Bronsted base.
hydrazine is used as a fuel in some rockets.
The most important hydride of phosphorus is phosphine, a gaseous analog of ammonia.
It is not possible to form phosphine by direct union of the elements.
There are two ways to prepare phosphine.
One method is by the action of an acid.
The compound burns in the air.
Phosphine is used as a fumigant for grains and in Semiconductor processing.
Like ammonia, phosphine and hydrogen halides form Phosphine compounds.
Phosphine is a weaker base than ammonia, so the compounds break down in water and escape from solution.
The offensive odor of rotten eggs is caused by hydrogen sulfide, a gas that is odorless.
Handling hydrogen sulfide is necessary because it is as toxic as hydrogen cyanide.
One doesn't smell hydrogen sulfide after a short exposure because it paralyzes the olfactory nerves.
The yield is low and the production of hydrogen sulfide by the direct reaction of elements is unsatisfactory.
A metal sulfide reaction with a dilute acid is a more effective method of preparation.
Unless a large amount of the oxidizing agent is present, the sulfur in H2S oxidizes to sulfur.
The oxidation of H2S in volcanic gases may be the cause of the deposits of sulfur.
Hydrosulfuric acid is formed by hydrogen sulfide dissolving in water.
The acid ionizes in two stages, yielding hydrogen sulfide ion in the first stage and hydrogen sulfide ion in the second stage.
The pure hydrogen halides are gases at room temperature.
The general techniques used to prepare other acids can be used to prepare the halides.
The bromine reacts with hydrogen to form hydrogen halide.
This reaction is important for preparing hydrogen bromide and hydrogen chloride.
A hydrogen halide can be produced by the acid-base reaction between a strong acid and a metal halide.
The reaction to completion is driven by the escape of the hydrogen halide.
Hydrogen chloride can be produced commercially and in the laboratory by the reaction of concentrated sulfuric acid with a chloride salt.
The least expensive chloride is usually sodium chloride.
Hydrogen bromide and hydrogen iodide can't be prepared with sulfuric acid because it's an oxidizer.
It is possible to prepare hydrogen bromide and hydrogen iodide with an acid such as phosphoric acid because it is a weaker oxidizing agent.
They are strong acids with the exception of hydrogen fluoride.
Salts of the halides can be produced by reactions of hydrohalic acids with metals, metal hydroxides, oxides, or carbonates.
chloride salts can be found in water.
The exceptions are AgCl, PbCl2, and Hg2Cl2.
It can frost or etch glass if it is attacked by hydrogen fluoride.
In production of hydrochlorofluorocarbons, hydrogen fluoride is used the most.
The manufacture of cryolite, Na3AlF6, is important in the production of aluminum.
The acid is important in the production of other inorganic fluorides, which serve as catalysts in the industrial synthesis of certain organic compounds.
It is relatively cheap to make hydrochloric acid.
It is important for the manufacture of metal chlorides, dyes, glue,glucose, and various other chemicals.
An acid used to remove oxide coating from iron or steel that is to be galvanized, tinned, or enameled is important for the activation of oil wells.
By comparison, the amounts of hydrobromic acid and hydroiodic acid used for commercial purposes are insignificant.
Most of the chemistry of carbon isn't relevant to this chapter.
The chapter covering organic chemistry will include other aspects of the chemistry of carbon.
The carbonate ion and related substances will be the focus of this chapter.
It is possible to prepare carbonates of the metals of groups 1 and 2 by the reaction of carbon dioxide with the respective oxide or hydroxide.
The carbonates form when you mix a solution of alkali metal carbonate with a solution of salts of the metals.
Tin(II) behaves differently in this reaction as carbon dioxide and the corresponding oxide form instead of carbonate.
Representative metal hydrogen carbonates such as NaHCO3 and CsHCO3 form when carbon dioxide is added to a solution of the hydroxides.
alkaline earth carbonates can be dissolved in water with carbon dioxide because hydrogen carbonate salts form.
stalagmites and stalactites are formed in caves when drops of water containing dissolved calcium hydrogen carbonate evaporate to leave a deposit of calcium carbonate.
The two carbonates used in the largest quantities are sodium carbonate and calcium carbonate.
The mineral trona, Na3(CO3)(HCO3)(H2O2), is found in the United States.
The reaction of calcium carbonate with hydrochloric acid is shown.
Hydrogen carbonates act as weak acids and weak bases.
Baking soda is a type of carbonate.
Baking soda and a solid acid are found in baking powder.
The majority of pure nitrogen comes from the fractional distillation of liquid air.
The atmosphere is made up of 70% nitrogen by volume.
There are more than 20 million tons of nitrogen on the earth's surface.
Nitrogen is a component of genes in plants and animals.
Nitrogen is a tasteless and odorless gas.
The temperature was 77 K when it boiled and 63 K when it froze.
Liquid nitrogen is inexpensive and has a low boiling point.
Nitrogen is unreactive because of the strong triple bond between the nitrogen atoms.
The only common reactions at room temperature are with hydrogen or oxygen and with Li3N, with certain transition metal complexes.
The ability of somebacteria to synthesise nitrogen compounds using atmospheric nitrogen gas as the source is one of the most exciting chemical events on our planet.
Nitrogen fixation is the process where organisms convert atmospheric nitrogen into biologically useful chemicals.
Nitrogen fixation occurs when lightning passes through air, causing nitrogen to react with oxygen to form nitrogen oxides, which are then carried down to the soil.
Nitrogen compounds are required for survival.
Most organisms can't absorb nitrogen from the atmosphere.
The strong nitrogen-nitrogen triple bond makes atmospheric nitrogen very unreactive.
A few organisms can overcome this problem through a process called nitrogen fixation.
A few organisms are able to process nitrogen.
Nitrogen fixation occurs when organisms convert atmospheric nitrogen into useful chemicals.
The only organisms capable of nitrogen fixation are microorganisms.
Iron and Molybdenum are contained in the nitrogenases that these organisms use.
This OpenStax book is available for free at http://cnx.org/content/col11760/1.9 and is the best-known example of a symbiotic relationship with plants.
The main starting material for making ammonia is large volumes of atmospheric nitrogen.
When a chemical process requires an atmosphere, it's helpful.
When canned foods and luncheon meats are sealed in nitrogen, they retain a better flavor and color, and are less likely to oxidize.
Fresh produce can be available year-round with this technology.
Nitrogen is present in all of the oxidation states.
Nitrogen can be reduced by some active metals.
We will look at nitrogen-oxygen chemistry in the rest of the section.
Nitrogen exhibits each of its positive oxidation numbers from 1+ to 5+ in well-characterized nitrogen oxides.
Water and nitrous oxide form when ammonium nitrate is heated.
Nitrogen gas, oxygen gas, and water vapor are generated by heating.
This reaction can be very dangerous and should never be attempted by anyone.
In 1947 there was a major explosion in Texas City, Texas, and in the same year there was another explosion in West, Texas.
Over the last 100 years, there have been nearly 30 similar disasters around the world, resulting in the loss of many lives.
The nitrogen in the nitrate ion oxidizes the nitrogen in the ammonium ion.
N2O has these structures in the left and right.
When lightning strikes in the air, no forms.
The method of preparing nitric oxide is burning ammonia.
The best way to prepare nitric oxide is in the laboratory.
It is one of the air pollutants created by internal combustion engines when atmospheric nitrogen and oxygen react.
nitric oxide is a gas at room temperature.
When a molecule with an unpaired electron combines with another molecule to form a bond, it's normal.
Liquid and solid NO both contain N2O2 dimers, like that shown in Most substances with unpaired electrons exhibit color by absorbing visible light; however, NO is odorless because the absorption of light is not in the visible region of the spectrum.
The equilibrium between NO and N2O2 is shown.
The liquid and solid states contain Dinitrogen trioxide.
It becomes a mixture of NO and NO2.
The N2O3 is only found in liquid or solid states and has two structures.
Nitrogen dioxide can be prepared by using air and nitric oxide.
Cu(NO3)2 and brown fumes of NO2 are produced by the reaction of copper metal with concentrated HNO3.
The nitrogen dioxide molecule has an unpaired electron which is responsible for its color and paramagnetism.
The dimerization of NO2 is also responsible for it.
Nitrogen dioxide has a deep brown color due to the presence of NO2 molecule.
The color is almost completely gone at low temperatures.
The structures of nitrogen dioxide and dinitrogen tetraoxide are shown.
The molecule of dinitrogen pentaoxide, N2O5 is shown in this image.
The oxides of nitrogen react with water and form oxyacids.
HNO2 forms when N2O3 reacts with water.
Nitrogen(IV) oxide, NO2, disproportionates in one of two ways when it reacts with water because there are no stable oxyacids containing nitrogen with an oxidation state of 4+.
HNO3 and NO will form at higher temperatures.
When heated with substances, nimbus oxide resembles oxygen.
N2O oxidizes when heated to form nitrogen and oxygen.
nitrous oxide supports combustion better than air because one-third of the gas liberated is oxygen.
A bottle of this gas has a glowing splinter in it.
Nitric oxide is both an oxidizer and a reducing agent.
A solid or burned form of P4O10 is formed when the phosphorus distills out of the furnace.
P4O10 is the beginning of the preparation of many other compounds.
In the chemical industry, the acids andphosphates are useful.
The manufacture of ferrophosphorus and phosphor bronze is one of the uses.
Pesticides, matches, and some plastic are made with Phosphorus.
There is a metal called Phosphorus.
In compounds, the oxidation states are 3-, 3+, and 5+.
The heads of strike-anywhere matches include P4S3.
There are two oxides of Phosphorus, P4O6 and P4O10, both shown in the picture.
The vapor is poisonous.
It oxidizes slowly in the air and forms P4O10 when heated to 70 degrees.
Phosphorous acid, H3PO3 is formed slowly in cold water.
The structures of P4O6 and P4O10 are shown in the image.
P4O10 is a white powder that is prepared by burning excess oxygen.
Its enthalpy is very high and it is a poor oxidizer.
The halogens will react directly with Phosphorus.
The trihalides are more stable than the nitrogen trihalides because of nitrogen's inability to form more than four bonds.
The liquid is made by passing chlorine over molten phosphorus.
The trichloride is oxidized with excess chlorine to make an off-white solid.
When heated, the trichloride and chlorine form an equilibrium with the pentachloride.
The image shows the structure of PCl3 and PCl5 in the gas phase.
Like most other nonmetal halides, PCl3 and PCl5 react with an excess of water and yield hydrogen chloride and phosphoric acid.
The compounds readily react with the Lewis bases to give the anion.
There are also the following compounds: 4, 4, 4, 4, 4, 4, 4, 4, 4, 4, 4, 4, 4, 4, 4, 4, 4,
Oxygen is abundant on the earth's surface.
The surface of the earth is made up of three parts.
Oxygen and other elements make up 50% of the mass of the earth's crust.
Oxygen occurs as O2 and O3 in the air.
It makes up 20% of the air.
Oxygen is a large part of plants and animals.
Oxygen is tasteless and odorless at normal temperatures.
It is denser than air.
Oxygen is very important to aquatic life because it is only slightlysoluble in water.
Oxygen comes from air and water.
The air is cooled and compressed before it liquefies.
Oxygen has a higher boiling point than nitrogen, which has a lower boiling point.
The other components of air can be separated based on their boiling points.
Oxygen is needed in the burning of fuels.
Plants and animals breathe in air.
The administration of oxygen-enriched air is an important medical practice when a patient is receiving an inadequate supply of oxygen.
Oxygen is used in the chemical industry to oxidize many substances.
Oxygen produced commercially is important in removing carbon from iron during steel production.
In metal fabrication and in the cutting and welding of metals with torches, large quantities of pure oxygen are required.
The space industry depends on liquid oxygen.
It's an oxidizer in rocket engines.
It is the main source of oxygen in space.
Oxygen is important to life.
Slow oxidation of chemical compounds provides the energy required for the maintenance of normal body functions in humans and other organisms.
Oxygen is the oxidizer in these reactions.
Oxygen enters the blood from the lungs, where it is combined with hemoglobin to make oxyhemoglobin.
This form of blood transfers oxygen to the tissues.
Carbon dioxide and water are the ultimate products.
The carbon dioxide is carried through the veins to the lungs, where it is collected by another supply of oxygen.
The energy liberated from the oxidation of the food is the same as if it were burned outside the body.
The oxygen that became carbon dioxide and water by the metabolism of plants and animals returns to the atmosphere.
Ozone contributes to the sharp odor associated with sparking electrical equipment.
The bent ozone molecule is shown in the image.
A layer of the atmosphere extending from 10 to 50 kilometers above the earth's surface is called the stratosphere.
Skin cancer can be prevented by the presence of stratospheric ozone.
It has been shown that chlorofluorocarbons, which were present as aerosol propellants in spray cans and as refrigerants, caused the decline in ozone.
This happened because ultraviolet light causes chlorine to break down.
The chlorine atoms react with ozone and remove O3 from the air.
This process is covered in detail in our coverage.
The ozone hole is starting to decrease in size as atmospheric concentrations of atomic chlorine decrease.
Ozone in the troposphere is a problem.
Ozone is a component of photochemical smog.
It can be used as a bleaching agent for oils, waxes, fabrics, and starch.
It is an alternative to chlorine.
A strong oxidizer is oxygen.
It reacts with many compounds.
There is a reaction of oxygen and iron.
Superoxides are formed by the more active metals.
Nonmetals give oxides.
Some compounds react with oxygen.
Further oxidation by oxygen can occur if it is possible to oxidize any elements in a compound.
H2S has sulfur with an oxidation state of 2-.
We would expect H2S to react with oxygen because the sulfur does not exhibit its maximum oxidation state.
It gives water and sulfur dioxide.
The difficulty of removing electrons from oxygen in most oxides is mirrored by the ease with which oxygen picks up electrons.
Oxygen can be formed from oxides of the elements by the very reactive fluorine.
All metals form oxides.
The oxides of most representative metals can be produced by heating the corresponding hydroxides or carbonates.
The oxidation-reduction reactions created by heating nitrates or hydroxides lead to the creation of alkali metal oxides.
The member for which the inert pair effect is most pronounced forms an oxide in which the oxidation state of the metal ion is two less than the group oxidation state.
Tl2O, PbO, and Bi2O3 form when burning thallium.
The lighter members of each group have oxides that show oxidation state.
It forms from burning tin.
When mercury is warm enough to form HgO, it forms slowly at higher temperatures.
Burning the members of groups 1 and 2 in the air is not a good way to form oxides of these elements.
These metals combine with nitrogen in the air to form oxides and ionic nitrides.
Several form superoxides when heated.
The oxide ion is a very powerful hydrogen ion acceptor.
The oxides of the representative metals react with acids to form salts.
In making firebrick, crucibles, furnace linings, and thermal insulation, magnesium oxide is important.
Blocks of calcium oxide heated by gas flames were the stage lights in theaters before electricity was available.
There are many uses for calcium oxide.
It emits a white light when it is heated.
Most of the useful products prepared from calcium oxide and calcium hydroxide do not contain calcium.
Representative metals, metalloids, and nonmetals are useful as bases.
The mineral corundum is a very hard substance used as an abrasive for grinding and polishing.
The jewelry trade depends on Corundum.
The wide variety of colors possible for sapphires is due to the presence of a small amount of chromium.
Artificial rubies and sapphires are made by melting aluminum oxide with small amounts of oxides to produce the desired colors and cooling the melt to produce large crystals.
Synthetic Ruby crystals are used in Ruby lasers.
Zinc oxide was a useful white paint color, but pollutants tend to oxidize it.
The compound is important in the manufacture of automobile tires and other rubber goods.
The zinc-oxide-based sunscreens are shown to help prevent sunburn.
The zinc oxide in these sunscreens is found in the form of small grains.
Lead dioxide is found in lead storage batteries.
Lead dioxide is a powerful oxidizer because it tends to reverting to the more stable lead ion by gaining two electrons.
The zinc oxide protects the exposed skin.
Strong oxidizers are important in chemical processes.
H2O2 is prepared from metal peroxides.
As the size of the cation increases, the stability of the peroxides and superoxides of the alkali metals increases.
The OH- ion is contained in Hydroxides.
Two general types of reactions can be used to prepare these compounds.
The metal or metal oxide can be reacted with water to producesoluble metal hydroxides.
The OpenStax book is free and can be found at http://cnx.org/content/col11760/1.9 hydroxide ion.
The metals of groups 1 and 2 react with water to form hydroxides and hydrogen gas.
The oxides form metal hydroxides.
Most other metal oxides are base anhydride and do not form hydroxides in water.
It is possible to prepare the insoluble hydroxides of magnesium, beryllium, and other representative metals by adding sodium hydroxide to a solution of a salt of the metal.
The Lewis acid-base reaction is the important aspect of complex ion formation in this chapter.
A cheap, strong base is used by the industry.
The oxide is the more expensive starting material for the production of NaOH.
In the United States, the top 10 chemicals in production are all sodium hydroxide.
There is a compound called sodium hydroxide.
40 grams of sodium hydroxide can be dissolved in 60 grams of water at 25 degrees.
In the production of other sodium compounds, as well as in the production of other chemicals, sodium hydroxide is used.
Many of the applications of hydroxides are for neutralization of acids, such as the antacid shown in and for the preparation of oxides by thermal decomposition.
The antacid milk of magnesia is constituted by the suspension of magnesium hydroxide.
Because of its ready availability, calcium hydroxide is used extensively in commercial applications that need a cheap, strong base.
The salts are prepared using the reaction of hydroxides with acids.
An antacid made from CaCO3 can be used to counteract the effects of acid in your stomach.
There is a link between chlorine and sodium hydroxide because there is an important process that makes both chemicals at the same time.
Many parts of the world have large deposits of sodium chloride, which is used in the chlor-alkali process.
The process oxidizes chloride ion to chlorine and produces sodium hydroxide.
The electrons travel through the outside electrical circuit.
Although the positive sodium ion migrates towards the negative electrode, metallic sodium doesn't form because it's too difficult to reduce under the conditions used.
Oxygen reacts with nonmetals to form oxides.
A variety of oxides might form depending on the available oxidation states for the element.
Oxygen and florine will combine to form OF2, a type of fluoride.
Sulfur dioxide, SO2, and sulfur trioxide are oxides of sulfur.
The smell of burning sulfur comes from sulfur dioxide.
The image shows the structure of sulfur dioxide.
Sulfur dioxide can be produced from either burning sulfur or roasting sulfides in the air.
Sulfurous acid, H2SO3 forms first, but quickly becomes sulfur dioxide and water.
When reducing agents react with hot, concentrated sulfuric acid, sulfur dioxide forms.
The SO2 molecule is bent when sulfur dioxide is at room temperature.
The structure of sulfur trioxide in the gas phase is shown in the image.
It is possible to prepare oxygen-halogen compounds by the reactions of the halogens with oxygen-containing compounds.
Oxygen compounds with chlorine, bromine, and iodine are oxides.
fluorine compounds with oxygen are fluorides.
The chemistry of the oxides is not important as a class.
The only commercially important compounds are dichlorine oxide and chlorine dioxide.
They are used for water treatment and as bleaching agents.
The structures of the (a) and (b) ClO2 molecule are shown in this image.
Acid anhydride are formed when nonmetal oxides are allowed to react with water.
oxyanions can form salts with metal ion.
N2O5 and NO2 react with water to form HNO3.
The separation of gold from silver was aided by the acid.
There are traces of nitric acid in the atmosphere after a storm.
Bengal saltpeter is found in India and other countries of the Far East.
The left and right images show the structure of nitric acid, HNO3 and its resonance forms.
It is often yellow or brown in color because NO2 forms as the acid degrades.
Concentrations of nitric acid are commercially available.
It is a strong oxidizer and a strong acid.
The action of nitric acid on a metal rarely produces H2 in large quantities.
The products are formed by the concentration of acid, the activity of the metal and the temperature.
Nitrogen dioxide is reduced by less active metals such as copper, silver, and lead.
The NO is produced by the reaction of copper and nitric acid.
In each case, the nitrate salts of the metals form.
It's useful in dissolving gold, Platinum, and other metals that are more difficult to oxidize than hydrogen.
One of the uses of nitric acid is to prepare metal nitrates.
In the manufacture of explosives, dyes, plastics, and drugs, it is important.
Salts of nitric acid can be used asfertilizers.
Gunpowder is a mixture of sulfur and charcoal.
A pale blue solution of nitrous acid, HNO2, is created by the reaction of N2O3 with water.
When it is heated, it disproportionates into nitric acid and nitric oxide.
A strong oxidizing agent oxidizes the acid to nitric acid.
This image shows the structure of a molecule.
Hot dogs and cold cuts have an added ingredient called NaNO2.
There are two functions of the nitrite ion.
It prolongs the red color of the meat and limits the growth ofbacteria that can cause food poisoning.
The addition of sodium nitrite to meat products is controversial because nitrous acid reacts with certain organic compounds to form a class of compounds known as nitrosamines.
Cancer can be produced in laboratory animals.
The FDA has limited the amount of NaNO2 in foods.
The acid can explode, but the nitrites are more stable.
AgNO2 is slightly soluble in water.
The common name of this compound is phosphoric acid, and is commercially available as a 82% solution known as syrupy phosphoric acid.
Many soft drinks have an added use for phosphoric acid.
When pure, H3PO4 has a Lewis structure.
Dilution of the products with water, followed by the removal of calcium sulfate, gives a solution contaminated with calcium dihydrogenphosphate, Ca(H2PO4)2, and other compounds.
It is possible to make pure orthophosphoric acid by dissolving P4O10 in water.
White crystals of phosphorous acid will appear when enough water is evaporated.
The odor of the crystals is similar to that of garlic.
Two hydrogen atoms are acidic in a molecule of phosphorous acid.
It is not possible to replace the third atom of hydrogen because it is not very acidic.
The preparation of sulfuric acid begins with the oxidation of sulfur to sulfur trioxide and then converting the trioxide to sulfuric acid.
Pure sulfuric acid is a liquid that is cold.
It fumes when it's heated because it's acid.
The heating process causes more sulfur trioxide to be lost than water.
A constant boiling solution and commercially concentrated H2SO4 is what the acid of this concentration is made of.
The amount of sulfuric acid used in industry is higher than any other compound.
There is a structure of sphuric acid.
The strong affinity of concentrated sulfuric acid for water makes it a good dehydrating agent.
It is possible to dry gases and immiscible liquids that don't react with acid.
The acid ionizes in two stages.
The first stage of the solution is complete.
Sulfates, such as Na2SO4 and hydrogen sulfates, are formed by sulfuric acid.
Sulfates of barium, strontium, calcium, and lead are slightly soluble in water.
The HSO4 ion is an acid, hydrogen sulfates, such as NaHSO4, exhibit acidic behavior, and this compound is the primary ingredient in some household cleanser.
An oxidizing agent is hot, concentrated sulfuric acid.
Sulfur dioxide is dissolved in water to form a solution of sulfurous acid.
Sulfurous acid can't be isolated because it is unstable.
The sulfur dioxide is expelled by heating a solution of sulfurous acid.
Sulfurous acid is strong.
It is necessary to add a base to a sulfurous acid solution and then evaporate the water in order to prepare solid sulfite and hydrogen sulfite salts.
The reaction of SO2 with oxides and hydroxides forms these salts.
Sulfates are always present in solutions of sulfites.
The hypohalous, halous, halic, and perhalic acids are represented by the compounds HXO, HXO2, and HXO3.
The perhalic acids are very strong and the hypohalous acids are weak.
Table 18.2 has a list of known acids and their pKa values.
There are no known salts for this compound.
It is not certain if the name hypofluorous acid is appropriate for HOF.
The reactions of chlorine and bromine with water are similar to those of fluorine with ice, but they don't go to completion and result in mixed hypohalous and hydrohalic acids.
Hypohalous acids are only found in solution.
The hypohalous acids are weak, but HOCl is a stronger acid than HOI.
Adding base to solutions of hypohalous acids produces salts with the basic hypohalite ion, OX-.
It is possible to separate these salts from the rest.
The reaction to hypochlorite is slow because of the unstable nature of the hypohalites.
The commercial preparation involves the lysis of cold, dilute, aqueous sodium chloride solutions under certain conditions where the resulting chlorine and hydroxide ion can react.
Chlorous acid is not stable, it slowly degrades in solution to produce chlorine dioxide, hydrochloric acid, and water.
The chlorite ion is a strong oxidizer and can be used in the bleaching of paper because it does not damage the paper.
chlorous acid reacts with bases to produce chlorineite ion.
Chloric acid, HClO3 and bromic acid are stable only in solution.
The reaction is similar to chlorous acid.
The halic acids have strong acids and active oxidizing agents.
Another preparative method is the oxidation of a hot solution of a metal to form the appropriate metal chlorates.
It is a weed killer and an oxidizing agent.
When halic acids react with bases, chlorineates are produced.
When treating a perchlorate, such as potassium perchlorate, with sulfuric acid under reduced pressure, perchloric acid forms.
Perchloric acid and its salts are powerful oxidizing agents as they are more stable in a lower oxidation state than 7+.
There have been serious explosions when heating concentrated solutions.
When perchloric acid is cold, its reactions are slow.
The acid is very strong.
The perchlorate ion is found in most salts.
It is possible to prepare them from the reactions of bases with perchloric acid and hot solutions of their chlorides.
When perchloric acid reacts with a base or with hot solutions of their chlorides, perchlorate ion can be produced.
The best per bromate salts involve the oxidation of bromates in basic solution with fluorine gas followed by acidification.
There are few uses for this acid or its salts.
Metaperiodic acid, HIO4 and paraperiodic acid are some of the different acids that have iodine in them.
The acids react with bases to form salts.
You will be able to describe the properties, preparation, and uses of sulfur Sulfur in nature by the end of this section.
H2S is contained in volcanic gases.
Representative metals, metalloids, and nonmetals are underground in Texas and Louisiana.
The underground deposit is caused by superheated water being forced down the outermost of three pipes.
The sulfur is melted by the hot water.
The liquid sulfur is compressed into the innermost pipe.
The air causes the liquid sulfur to flow through the outlet pipe.
Transferring the mixture to large vats allows the sulfur to separate.
This sulfur is pure and does not require purification for most uses.
The Frasch process is used to mine sulfur.
Hydrogen sulfide comes from the purification of natural gas.
There are several allotropic forms of sulfur.
The true formula is S8 because of the eight-membered rings in the stable form.
We will follow the practice of chemists who use S to simplify coefficients in chemical equations.
sulfur is a member of group 16 and exhibits a nonmetallic behavior.
It oxidizes metals and is available for free in the OpenStax book at http://cnx.org/content/col11760/1.9.
Sulfur oxidizes less nonmetals, and more nonmetals, such as oxygen and the halogens, will oxidize it.
Other oxidizers oxidize sulfur.
Sulfur forms many compounds in which it exhibits positive oxidation states.
The elements in group 17 are the halogens.
The elements are fluorine, chlorine, bromine, and astatine.
The elements are not free to occur in nature.
The occurrence, preparation, and properties of halogens will be examined in this section.
Next, we will look at the interhalogens and the representative metals.
The section will end with applications of halogens.
All of the halogens are in the water.
CaF2, Ca(PO4)3F, and Na3AlF6 are minerals.
In the Great Salt Lake and the Dead Sea, there are extensive salt beds that contain NaCl, KCl, or Magnesium.
hydrochloric acid is a component of stomach acid and part of the chlorine in your body.
The Dead Sea has underground brines.
Small amounts of Iodine compounds are found in the saltpeter, underground brines, and sea kelp.
Iodine is needed for the function of the thyroid gland.
The best sources of halogens are halide salts.
Depending on the ease of oxidation of the halide ion, it is possible to oxidize it.
iodide is the easiest oxidizer.
The main method of preparing fluorine is electrolytic oxidation.
A molten mixture of hydrogen fluoride, KHF2, and anhydrous hydrogen fluoride is the most common method of electrolysis.
lysis causes fluorine gas to form at the anode and hydrogen to form at the cathode.
The separation of the two gases is needed to prevent the formation of hydrogen fluoride.
The majority of commercial chlorine comes from the chlor-alkali process, in which the chloride ion in the solution of sodium chloride is electrolysis.
Chlorine is a product of the production of metals from their fused chlorides.
The chlorine can be prepared by the oxidation of the chloride ion in acid solution with strong oxidizing agents.
Representative Metals, Metalloids, and Nonmetals Chlorine is a stronger oxidizer than bromine.
This method is important for the production of domestic bromine.
The oxidation of iodic acid, HlO3 makes up some of the iodine.
Fluorine is a pale yellow gas, chlorine is a greenish-yellow gas, bromine is a deep reddish-brown liquid, and iodine is a grayish-black solid.
Liquid bromine has a high Vapor Pressure, and the reddish Vapor is visible in Iodine crystals.
These crystals form a beautiful deep violet Vapor when gently heated.
chlorine is a pale yellow-green gas, bromine is deep orange, and iodine is purple.
violet solutions of I2 molecule can be found in chloroform, carbon tetrachloride, carbon disulfide, and many hydrocarbons.
Brown solutions were given by the slightly dissolved iodine.
It forms brown solutions in the solutions of iodides.
The florine oxidizes an element to its highest oxidation state.
chlorine gives bromine and scl2.
Iodine doesn't react with sulfur.
The most powerful oxidizer of the elements is fluoroine.
The reverse reaction, the oxidation of fluorides, is very difficult because it spontaneously oxidizes most other elements.
The lighter noble gases (He, Ne, and Ar) are not affected by the reaction.
Many substances ignite on contact with florine, it is a strong oxidizing agent.
The OpenStax book is available for free at http://cnx.org/content/col11760/1.9 fluorine.
The adherent film of the fluoride salt makes it possible to handle fluorine in copper, iron, or nickel containers.
There is only one element that reacts with noble xenon gas.
Chlorine is less active than fluorine.
The reaction between chlorine and hydrogen in the dark is very slow.
Exposure of the mixture to light causes it to explode.
Chlorine is less active towards metals than fluorine.
Most nonmetals are attacked by chlorine, forming covalent compounds.
Adding multiple bonds or substitution is how chlorine reacts with compounds that only contain carbon and hydrogen.
This disproportionation is incomplete, so chlorine water is an equilibrium mixture.
The nonmetal chlorine is more negative than any other element.
We would expect chlorine to oxidize all of the other elements except for the three that are nonreactive noble gases.
The oxidizing property is responsible for its main use.
The reactivity of bromine is less than that of chlorine, but it has the same chemical properties.
The most reactive of the halogens is Iodine.
The weakest oxidizer is the iodide ion.
heating is required when iodine reacts with metals.
It doesn't oxidize other halides.
Compared with the other halogens, the reaction with water is slightly different.
A deep blue color is formed by a mixture of the two substances in the water.
The reaction is very sensitive for the presence of iodine in water.
The representative metals have been prepared.
The salts are important.
A salt is an ionic compound composed of cations and anions.
The salts can be prepared from metals, oxides, hydroxides, or carbonates.
We will show the reactions used to prepare salts.
The majority of the halides are ionic.
The halide of the metal is produced by the direct reaction of a metal and a halogen.
The reactions of the alkali metals are violent.
They give exciting demonstrations for budding students of chemistry.
There is a picture of the surface that is exposed to the residual mineral oil and the removal of the coating of sodium hydroxide.
The reaction with chlorine gas goes very well.
If a metal has two oxidation states, it may be necessary to control the stoichiometry in order to get the halide with the lower oxidation state.
There is a chapter on the stoichiometry of chemical reactions.
Solid HgI2 forms when solutions of Hg(NO3)2 are mixed.
There are many halides in nature.
There are many halides in the ocean and underground brines.
The source of magnesium ion used in the production of magnesium is found in the ocean.
Most of the compounds containing these elements come from these deposits.
One example is the chlor-alkali process.
Underground deposits of sodium chloride can be found all over the world.
There is a tunnel in the salt mine.
Interhalogen molecule consist of one atom of heavier halogen bonded by single bonds to an odd number of lighter halogen atoms.
There are three structures in IF3: T-shaped, square pyramidal, and pentagonal bipyramidal.
fluorine is able to oxidize iodine to its maximum oxidation state, 7+, whereas bromine and chlorine are more difficult to oxidize, achieve only the 5+-oxidation state.
The limit for the halogens is a 7+oxidation state.
The maximum number of smaller atoms can be increased as the radius of the larger atom increases.
Many of these compounds are unstable.
The interhalogens are similar to their components in that they are stronger oxidizing agents.
Interhalogens are closely related to the ionic polyhalides of the alkali metals.
There are many uses for the fluoride ion and fluorine compounds.
The compounds of carbon, hydrogen, and fluorine are being used as refrigerants.
The material Teflon is composed of units.
Some toothpastes and water supplies contain floride ion to fight tooth decay.
The teeth from Ca5(PO4)3(OH) are partially converted into Ca5(PO4)3F.
Representative metals, metalloids, and nonmetals chlorineine are important to bleach wood and cotton cloth.
Hypochlorous acid is formed when chlorine reacts with water.
In the production of compounds such as chloroform and ethyl chloride, large quantities of chlorine are important.
It is important to kill thebacteria in community water supplies with chlorine.
The production of certain dyes and the use of bromides as sedatives is important.
Light-sensitive silver bromide was a part of photographic film.
It is antiseptic to have Iodine in alcohol solution.
Iodide salts are needed for the proper functioning of the thyroid gland, and a deficiency may lead to the development of a goiter.
Iodized table salt contains a small amount of a radioactive substance.
It was important in the production of photographic film that silver iodide be used in the seeding of clouds to induce rain.
The elements in group 18 are noble gases.
They earned the name "noble" because they were assumed to be nonreactive.
The assumption was proved to be false by Dr. Neil Bartlett at the University of British Columbia.
The elements are present in a small amount.
Natural gas has a small amount of helium by mass.
Helium is isolated from natural gas by liquefying the condensable components.
Most of the world's commercial supply of this element can be found in the United States.
Liquid air is fractionally distilled from it.
There are other radioactive elements.
The radioactive gas is present in very small amounts in soils and minerals.
Lung cancer is a health hazard due to itsAccumulation in well-insulated, tightly sealed buildings.
The boiling points and melting points of noble gases are very low compared to other substances.
Weak London dispersion forces can only hold the atoms together when the motion of the atoms is very slight, as it is at very low temperatures.
Helium does not solidify on cooling at normal pressure.
It is liquid close to absolute zero at ordinary pressures.
Helium is safer to use than hydrogen because it does not burn.
Nitrogen is a narcotic, but helium is not.
Divers working under high pressures need to have a mixture of oxygen and helium.
Nitrogen narcosis, the so-called rapture of the deep, can be avoided using a helium-oxygen mixture.
The melting and welding of easily oxidizable metals and many chemical processes that are sensitive to air can be done with helium.
It is essential for achieving the low temperatures necessary to produce superconduction in traditional materials used in powerful magnets and other devices to have liquid helium.
Magnetic resonance images are a common medical diagnostic procedure.
Liquid nitrogen is cheaper than the other common coolant.
Neon is used in lamps and signs.
The red glow of neon is caused by an electric spark passing through a tube of neon.
It is possible to change the color of the light by mixing mercury or argon with the neon or using glass tubes of a special color.
In the manufacture of gas-filled electric light bulbs, it was preferable to use nitrogen because of its lower heat conductivity and chemical inertness.
A mixture of mercury and argon is found in fluorescent tubes.
The third most abundant gas is rosin.
The flash tubes are used to take pictures.
KrF2 is unstable at room temperature.
Stable compounds of xenon form.
XeF2 forms after heating an excess of xenon gas and cooling it.
The material is stable at room temperature in a dry environment.
XeF6, XeF4 and XeF6 are all prepared in the same way, with the same amount of fluorine and fluorine excess.
Oxygen is used to replace fluorine atoms in the xenon fluorides.
Evidence of RnF2 comes from radiochemical techniques.
Stable compounds of helium and neon are not known.
The periodicity of the representative elements is the focus of this section.
The elements are in groups 1 to 2.
The elements are representative of metalloids and nonmetals.
The group 1 alkali metals are very reactive, readily form ions with a charge of 1+ to form ionic compounds that are usuallysoluble in water, and react vigorously with water to form hydrogen gas and a basic solution of the metal hydroxide.
The outer electron of the alkali metals is more difficult to remove than the outer electron of the alkaline earth metals.
The metals exhibit an oxidation state of 2.
Hydrogen is more difficult to oxidize than aluminum, indium, and thallium.
thallium also occurs as the Tl+ ion in the oxidation state of aluminum, gallium, and indium.
Stable divalent cations and covalent compounds are formed by tin and lead.
Reduction from naturally occurring compounds is needed to produce the representative metals in their pure forms.
In the production of aluminum, lysis is important.
The main method for isolating magnesium, zinc, and tin is chemical reduction.
The other representative metals have similar procedures.
The metals are separated from the nonmetals in the periodic table by the elements.
The properties of metalloids and semimetals are similar to those of metals.
The structures of these elements are similar to those of nonmetals, but they are electrical.
Nonmetals have structures that are different from metals because they have more electrons that are bound to individual atoms.
Acid anhydride oxides react with water to form acidic solutions.
Many of the nonmetals have multiple allotropes that have different physical properties.
The chemistry of hydrogen is unique and it is the most abundant element in the universe.
Hydrogen has many of the same chemical properties as 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 nonmetals.
It reacts with oxygen, fluorine, and chlorine, less readily with bromine, and less readily with iodine, sulfur, and nitrogen.
The oxides of metals with lower reduction potentials are formed by hydrogen.
When dissolved in water, the hydrogen halides are acidic.
The preparation of the carbonates of the alkali and alkaline earth metals is usually done by reaction of an oxide or hydroxide with carbon dioxide.
Other carbonates are formed by precipitation.
Limestone, antacid Tums, and baking soda are examples of metal carbonates.
Most of the carbonates and hydrogen carbonates break down on heating.
Nitrogen has oxidation states ranging from 3 to 5.
Nitrogen can be reduced by the alkali metals and alkaline earth metals.
Nitrogen oxides and nitrogen hydrides are important substances.
Group 15 has oxidation states of 3- with active metals and of 3+ and 5+ with more nonmetals.
Oxygen and halogens oxidize phosphorus.
The oxides are P4O10 and P4O6.
There are two ways to prepare orthophosphoric acid, H3PO4, Chapter 18 | Representative Metals, Metalloids, and Nonmetals.
Three types of salts are formed by triprotic acid.
Oxygen is a reactive element.
The chemistry of oxygen is rich and well understood because of its reactivity and abundance.
The most common method for producing oxides is heating the corresponding hydroxides, nitrates, or carbonates.
The formation of superoxides can be caused by heating the metal or metal oxide.
The oxides form hydroxides by dissolving in water.
Most metals oxides react with acids.
The representative metals react with acids in acid-base reactions to form salts and water.
There are many uses for the hydroxides.
Most of the nonmetal oxides are acid anhydride-based.
The hydrogen atoms have to bond to the oxygen atoms in the molecule rather than the nonmetal atoms.
The strength of the oxyacid increases with the number of oxygen atoms in the nonmetal atom.
Sulfur forms the sulfide ion, S2, when it reacts with almost all metals.
Most nonmetals react with sulfur.
The halogens form halides.
The halides of the metals are different from the nonmetals.
Representative metal halides can be produced with solutions of the hydrohalic acids or directly with the elements.
Adding hydrohalic acids to compounds that contain basic anions is one of the laboratory preparations.
The most important property of the noble gases is their lack of activity.
They happen in low concentrations.
They find uses for atmospheres, neon signs, and coolants.
The three heaviest noble gases form fluorides.
The starting materials for a few other noble gas compounds are the xenon fluorides.
Predict the formulas for the nine compounds that may form when each species in column 1 reacts with each species in column 2.
Both strontium chloride and sodium chloride are white.
Slaked lime is used in the construction industry to make mortar and plaster.
The elements are in the same place.
The bonds in PbCl2 are different from the bonds in PbCl4.
Balance the equations for the reactions occurring at the electrodes and the overall reaction in the electrolysis of molten lithium chloride.
You may want to read the chapter on chemistry.
Chlorine can be produced by the electrolysis of molten sodium chloride.
In each case, calculate the mass of chlorine produced.
You may want to read the chapter on chemistry.
Magnesium is an active metal that burns in the form of powder, ribbons, and filaments to provide flashes of light.
A sample of a type of metal (an alloy of Sn, Pb, Sb, and Cu) is dissolved in nitric acid and metastannic acid is precipitated.
She leaves a small amount of tin(IV) oxide when she drives off the water.
Assume a 100% yield for the production of 100 kilograms of sodium metal using a current of 50,000 A.
Give the hybridization of the metalloid and the geometry of the compounds.
You can review the chapters on chemical bonding and advanced covalent bonding.
Write a Lewis structure for each molecule.
There is a chapter on chemical bonding.
The complete electron configuration for Silicon should be written using only the periodic table.
You can review the chapter on electronic structure.
There is a chapter on chemical bonding.
Silicon reacts with sulfur.
Determine the empirical formula of Silicon sulfide if 0.0923 g of Silicon reacts with sulfur.
A hydride of Silicon prepared by the reaction of Mg2Si with acid exerts a pressure of 306 torr in a bulb with a volume of 57.0 mL.
Imagine you found a diamond encased in a rock.
Two of the allotropes are made of carbon.
There is a diamond structure in Silicon.
Nitrogen is a very stable diatomic molecule.
In order to increase electronegativity, the following should be arranged.
The oxidation-reduction reaction is also a part of the reaction.
The changes in oxidation number that occur in the reaction are identified.
Lewis structures show that a hydrogen atom forms only one bond in a compound.
1.720 g of anhydrous Na2CO3 is left after the removal of the water of hydration.
The central nitrogen and the nitrogen atom should be indicated.
Determine the oxidation state of nitrogen.
There is a chapter on chemical bonding.
Predict the ONO bond angle and draw the Lewis structure for each of the following.
The chapters on chemical bonding and advanced theories of covalent bonding may be of interest to you.
HNO3 is 35.27 liters.
There is a chapter on chemical bonding.
Only one proton of the acid molecule reacts when H3PO3 is titrated.
Write equations that show the stepwise ionization of phosphorous acid.
One of the acids used in some cola drinks is phosphoric acid.
The Phosphorus(V) oxide is prepared by burning it.
The product of burning air is predicted.
Use equations to describe the reaction of water and potassium oxide.
Write a balanced chemical equation for the reaction of an excess of oxygen.
Oxygen is a strong oxidizer and tends to oxidize an element to its maximum oxidation state.
You can review the chapter on acid-base equilibria.
In SO2, SO3 and H2SO4 give the oxidation and hybridization state for sulfur.
Determine the oxidation state of sulfur.
Single bonds in S8 are formed by sulfur and oxygen.
Write two balanced chemical equations in which sulfuric acid acts as an oxidizer.
The reactions are similar to the industrial chemicals.
Lewis structures can be drawn for each of the following.
There is a chapter on chemical bonding.
The solution for contact lens is prepared to match the concentration in the body.
The hybridization of xenon should be given in each of the following.
The chapter on advanced theories of covalent bonding may be of interest to you.
There is a chapter on chemical bonding.
There is a chapter on chemical bonding.
A mixture was heated.
A sample of the white solid that formed reacted with hydrogen to yield 81 mL of xenon and hydrogen fluoride, which was collected in water, gave a solution of hydrofluoric acid.
Write balanced chemical equations for the reactions involving xenon and determine the empirical formula for the white solid.
The basic solutions of Na4XeO6 are powerful oxidants.
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