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    General Chemistry IIExam III Review: Chapter 20

    CHAPTER 20

    Nuclear Chemistry

    1. Radioactivityi. Alpha rays bend away from positive plate toward a negative plate

    1. Positive chargeii. Beta rays bend towards positive plate

    1. Negative chargeiii. Gamma rays are unaffected by magnetic and electric fields

    1. Similar to X-rays, but shorter wavelengthb. Nuclear Equations

    1. Symbolic representation of a nuclear reaction

    ii. Positron mass = electron mass, but with opposite charge (+)iii. Total charge and total # of nucleons is conserved

    1. The sum of subscripts of reactants equals sum of subscripts of products

    2. The sum of superscripts of reactants equals sum of superscripts of products

    c. Nuclear Stabilityi. Nuclear force strong attraction between nucleons

    1. Only effective at distances of 10 -15 cm2. Protons farther apart will repel each other because of their

    electric charges; inside the nucleus, the protons are closeenough for the nuclear force to keep them together

    a. stable nucleusii. Shell Model of the Nucleus

    1. A nuclear model in which protons and neutrons exist in levels,or shells, analogous to the shell structure for electrons

    2. Magic number a number of nuclear particles that make anucleus very stable

    a. A magic number is the number of nuclear particles in acompleted shell of protons or neutrons

    b. For protons: 2, 8, 20, 28, 50, 82, 114c. For neutrons: 2, 8, 20, 28, 50, 82, 126

    3. Evidence for shell model:a. Helium nuclei (seen in alpha particle emission) are

    especially stable, and they contain 2 protons and 2neutrons (both are magic numbers)

    b. Uranium-238 decays to the final product 206/82Pbc. Other radioactive decays series end at 208/82Nd. Isotopes of all elements tend to favor arrangements

    with even numbers of neutrons and protons

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    4. Band of Stability a. For nuclides up to Z = 20, the ratio of protons to

    neutrons is about 1.1:1.0b. As Z increases, the ratio increases to 1.5:1.0

    i. Larger numbers of protons increasingly repulse

    each other in the nucleus, so more neutrons arerequired to give attractive nuclear forces tooffset repulsions

    c. No stable nuclei exist above Z = 83d. All elements with Z less than or equal to 83 have at least

    one stable nuclide, except technetium (Z = 43) andpromethium (Z = 61)

    d. Types of Radioactive Decayi. Alpha emission

    1. Emission of a helium nucleus from an unstable nucleus2. The product has an atomic number that is two less, and a mass

    number that is four less, than original nucleusii. Beta emission

    1. Emission equivalent to the conversion of a neutron to a proton2. The product has a number that is one more than that of the

    original nuclear, and the mass number remains the same3. Emission of a high-speed electron from an unstable nucleus

    iii. Positron emission1. Positron is a particle identical to an electron in mass, but has a

    positive charge instead of a negative charge2. Equivalent to the conversion of a proton to a neutron3. The product nucleus has an atomic number that is one less

    than that of the original nucleus, and the mass number remainsthe same

    iv. Electron capture1. The decay of a unstable nucleus by capturing/picking up an

    electron from an inner orbital of an atom2. In effect, a proton is change to a neutron, as in positron

    emission3. The product nucleus has an atomic number one less than the

    original nucleus, and the mass number remains the same4. When another orbital electron fills the vacancy in the inner-

    shell orbital created by electron capture, an x-ray photon isemitted

    5. Same effect as positron emission a. Rather than emitting a positron, it captures an electron,

    but both effectively reduce atomic number by 1 without changing the mass number

    v. Gamma emission1. Emission from an excited nucleus of a gamma photon,

    corresponding to radiation with a wavelength 10 -12

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    2. Decayed product nucleus are often unstable, so they undergogamma ray emission to get to a lower-energy state

    3. Often occurs very quickly after radioactive decay4. Metastable nucleus a nucleus in an excited state with a

    lifetime of at least one nanosecond (10 -9)

    a. In time, metastable nucleus decays by gamma emissionb. There is no change in atomic numbervi. Spontaneous fission

    1. The spontaneous decay of an unstable nucleus in which aheavy nucleus of mass number greater than 89 splits intolighter nuclei and energy is released

    e. Predicting which type of radioactive decayi. Nuclides to the left of the band of stability have a higher

    neutron:proton ratio than needed for stability1. These nuclides decay by beta emission2. In beta emission, a neutron becomes a proton

    ii. Nuclides to the right of the band have a lower neutron:proton ratiothan needed for stability

    1. These nuclides decay by positron emission or electron capture2. Both these processes convert a proton to a neutron

    iii. For example:1.

    6

    11 C

    5

    11 B +

    1

    0 e a. Positron emissionb. Neutron:proton ratio increases when Carbon-11 decays

    to boron-112.

    6

    14 C

    7

    14 N +

    1

    0e

    a. Beta emissionb. Neutron:proton ratio decreases

    3. Phosphorous-31 has atomic mass of 31.0amu; you can expect that Phorphorous-31 is a stable isotope. Phorphorous-30should be expected to undergo positron emission or electroncapture (positron emission is observed)

    iv. Positron emission and electron capture are competitive decayprocesses

    1. Rate of electron capture increases with atomic number of thenuclide, so it is more important in heavier elements

    2. Positron emission typically seen in lighter elements (like

    phosphorous-30)v. In very heavy elements (Z 83) alpha emission is common form of decay (uranium-238)

    2. Radioactive decay seriesi. All nuclides with Z > 83 are radioactive

    b. Radioactive decay series a sequence in which one radioactive nucleusdecays to a second, followed by a third, fourth, and so on

    i. Eventually, a stable nucleus, which is an isotope of lead, is reached

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    ii. There are three radioactive decay series found in nature:1.

    92

    238 U

    82

    206 Pb 2.

    92

    238 U

    82

    207 Pb 3.

    90

    232 Th

    82

    208 Pb

    Nuclear bombardment reactions3. Transmutation the change of one element to another by bombarding the nucleus of

    the element with nuclear particles or nucleia. Ernest Rutherford found that when alpha particles emitted from a

    radioactive element collided with a stable nitrogen nuclei, a protons wereejected in the process of forming an oxygen nucleus

    b. Bombardment of beryllium with alpha particles led to the discovery of theneutron in 1932, as the radiation from the beryllium atom was found to bewithout charge

    c. Abbreviated form for bombardment notationi.

    7

    14 N +

    2

    4 He

    8

    17 O +

    1

    1 H ii.

    7

    14 N ( , p)

    8

    17 O Particle AbbreviationNeutron nProton p

    Deuteron dAlpha

    d. Very heavy elements cant be transmutated by natural alpha emissions e. Particle accelerator necessary to shoot charged particles into heavy nuclei

    i. Device used to accelerate electrons, protons, alpha particles, and otherions to very high speeds

    ii. Kinetic energy of these particles is measured in electron volts (eV)1. An electron volt is the quantity of energy that would have to be

    imparted to an electron to accelerate it by on volt potential2. Particle accelerators typically give particles energies of

    millions of electron voltsiii. Cyclotron

    1. Type of particle accelerator with two hollow, semicircularmetal electrodes called dees, in which charged particlesaccelerate by stages to higher kinetic energies

    4. Transuranium elementsi. Elements with atomic numbers great than that of uranium (Z = 92)

    ii. Prepared in a laboratory by firing various particles at uraniumisotopes

    1. Plutonium-239 is prepared in large quantities in nuclearreactors, and is used for nuclear weapons

    2. Plutonium was first produced by accelerating a deuteron at auranium atom to yield neptunium-238, which decayed toplutonium-238 (beta decay)

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    iii. Glenn Seaborg proposed that the next elements to be discoveredbeyond actinium (Z=89), called actinides, should b e placed at thebottom of the periodic table under the lanthanides

    5. Radiation Countersa. Geiger counter ionization counter used to count particles emitted by

    radioactive nuclei, consisting of a metal tube filled with gas, like argoni. Radiation enters through thin glass or plastic windowii. A wire runs from tube center, from which wire is insulated

    iii. The tube and wire are connected to high-voltage source so that thetube becomes negative electrode, and the wire becomes positiveelectrode

    iv. Normally, gas in tube is insulator and no current flows, but whenradiation, like an alpha particle, passes through the window of thetube and into the gas, atoms are ionized

    v. Free electrons accelerate towards positively charged wirevi. These electrons collide with more atoms to set additional electrons

    freevii. This process gives a pulse of current, amplified as a click

    viii. Alpha and beta particles can be detected by a Geiger counter1. Neutrons can be detected only with boron trifluoride is added

    to the gas in the tube2. Neutrons react with boron-10 nuclei to produce alpha

    particles, which can then be detectedb. Scintillation counter

    i. A device that detects nuclear radiation from flashes of light generatedin material by the radiation

    ii. A phosphor is a substance that emits a flash of let when radiationstrikes it

    iii. The activity of a radioactive source is the number of nucleardisintegrations per unit time occurring in a radioactive material

    1. A curie (Ci) a unit of activity is equal to 3.7 x 10 10

    Rate of Radioactive Decay6. Rate of decay and half-life

    a. Rate = kN t i. k = radioactive decay constant (different for each nuclide)

    ii. N t = number of radioactive nuclei at time t b. The decay constant comes from the number of nuclear disintegrations over a

    period of timec. Given activity in curies (Ci), you can calculate rate:

    i. Rate = (activity in Ci) x [(3.7 x 10 10 nuclei/s)/(1.0 Ci)]ii. k = rate/ N t

    d. Half-life i. The time it takes for one half of the nuclei in a sample to decay

    ii. Related to radioactive decay:1. t 1/2 = 0.693/ k

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    2. k = decay constant iii. With decay constant, you can calculate the fraction of the radioactive

    nuclei that remain after any given time1. ln(N t /N 0) = -kt 2. If given half life, time elapsed, but no k :

    a. Plug in (-0.693t/t 1/2 ) for k b. Solve for the ratio of N t / N 0 7. Radioactive dating

    a. Carbon datingi. Carbon-14 has half life of about 5700 years

    ii. Present in atmosphere because of cosmic raysiii. A nitrogen-14 atom collides with a neutron from space, yielding a

    proton and a carbon-14 atomiv. By measuring the level of beta emissions in a dead object (or, the

    conversion of carbon-14 back into nitrogen-14) you can determinehow long it has been dead

    Mass-Energy Calculations8. Mass-energy equivalence

    a. When objects lose energy, they lose an equivalent amount of massb. E = mc 2 c. c = 3.00 x 10 8 m/s 2 d. Miniscule amount of mass lost as energy is lost in chemical reactionse. Nuclear reactions involve 1,000,000 times bigger mass changes per mole of

    reactant than chemical reactions9. Nuclear binding energy

    a. Binding energy the energy needed to break a nucleus into its individualprotons and neutrons

    b. Mass defect the nucleon mass minus the atomic massi. A larger positive difference indicates a more stable atom, as more

    energy is released when the nucleus formsii. Larger mass defect = more stable nucleus

    c. Nuclear fission a nuclear reaction in which a heavy nucleus splits intolighter nuclei and energy is released

    d. Nuclear fusion a nuclear reaction in which light nuclei combine to give amore stable, heavier nucleus plus possible several neutrons, and energy isreleased

    Nuclear Fission and Fusione. Both processes are ways for atoms of extreme size to attain a size closer to

    the intermediate range (mass number = 50)10. Nuclear fission and nuclear reactors

    a. When uranium-235 nucleus splits, about two or three neutrons are released,which are then absorbed by other uranium-235 nuclei, which then split andrelease more neutrons

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    b. Chain reaction a self-sustaining series of nuclear fissions cause by theabsorption of neutrons released form previous nuclear fissions

    c. Critical mass the smallest mass of fissionable material in which a chainreaction can be sustained

    i. If mass is too small, the neutrons released when a nucleus splits will

    be lost too quickly for the chain reaction to be sustainedii. Supercritical mass a mass much larger than the critical mass whichresults in huge numbers of rapidly splitting nuclei

    1. Atomic bombs contain two subcritical masses of uranium that,when pushed together with explosives, form a supercritical mass

    d. Nuclear fission reactori. A device that permits a controlled chain reaction of nuclear fissions

    ii. Fuel rods cylinders that contain fissionable materialiii. Control rods cylinders composed of substance that absorb neutrons,

    like boron and cadmium, to slow the cahin reaction

    iv. Moderator a substance that slows down neutrons, required if uranium-235 is the fuel

    1. Slows down nuclei released by uranium- 235 so that they arent absorbed by uranium-238, but by other uranium-235 atoms

    2. Heavy water, light water, and graphitee. Nuclear fusion

    i. Energy can be obtained by combining light nuclei into a heavynucleus, because of the bonding energy of the nucleus

    ii. To get two light nuclei to fuse, one nucleus must bombard enoughwith enough kinetic energy to overcome the repulsion of electriccharges

    iii. Cold fusion of deuterium atoms is best power source, but it unlikely iv. Much more feasible is hot fusion

    1. Tritium and deuterium are heated to 1 million degrees Celsius,at which point they form a plasma

    2. Plasma an electrically neutral gas of ions and electrons3. At this temperature, the plasma is essentially separate nuclei

    and electrons4. Hydrogen bombs work by first using a fission reaction to

    create the extreme temperatures needed for fusioncomponents in the bomb to react

    v. Plasma quickly loses heat when it touches any material1. Magnetic or laser fusion reactors would solve this problem

    CHAPTER 21

    1. General Observations About the Main-Group Elementsa. Elements on the left side- largely metallicb. Right side- largely nonmetallic

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    c. Metallic elements:i. Low ionization energies and low electronegativities

    ii. Tend to lose their valence elections to form cationsiii. Oxides of metals are usually basiciv. Oxides of most reactive metals react with water to give basic solutions

    v. Generally have oxidation states equal to the group number1. Some in fifth and sixth periods have ox states equal to thegroup number minus 2

    d. Nonmetallic elements:i. Form monoatomic anions and oxoanions

    ii. Oxides are acidiciii. Have a variety of oxidation states ranging from the group number

    (most positive value) to the group number minus eight (most negative value)

    e. Metallic characteristics of the main group elements in the periodic tablegenerally decrease in going across a period from left to right

    f. Metallic characteristics of the main group elements in the periodic tablebecome more important going down any column (group)

    g. Second period element is often rather different from the remaining elementsin its group

    i. Has a small atom that tends to hold electrons strongly highelectronegativity

    ii. Bonding only includes s and p orbitals, other elements may use dorbitals

    2. Metals: Characteristics and Productiona. Metal- lustrous, high electrical and heat conductivities, and is malleable and

    ductileb. Alloy- material with metallic properties that is either a compound or a

    mixture (can be homo- or heterogeneous)c. Sources of metals

    i. Mineral- naturally occurring inorganic solid substance or solidsolution with a definite crystalline structure

    ii. Ore- a rock or mineral from which a metal or nonmetal can beeconomically produced

    d. Metallurgy- scientific study of the production of metals from their ores andthe making of alloys having various useful properties

    i. Steps in the production of metal from its ore: preliminary treatment,reduction, refining

    e. Preliminary Treatment i. Separate mineral from ore

    f. Reductioni. If a metal isnt already free, must be obtained from one of its

    compounds by reduction (using electrolysis or chemical reduction)1. Electrolysis: uses an electric current to reduce a metal

    compound to the metal (lithium from lithium chloride)

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    2. Reduction: cheapest chemical reducing agent is some form of carbon such as hard coal or coke.

    g. Refiningi. Metal contains impurities and needs to be purified (refined) various

    techniques are used

    3. Bonding in Metalsa. Special properties of metals result from delocalized bonding (bondingelectrons are spread over a number of atoms)

    b. Electron-sea mo del: array of positive ions surrounded by a sea of valenceelections free to move over the entire metal crystal

    i. When metal is connected to electric current, electrons easily moveaway from negative side of electric source and toward positive side ELECTRIC CURRENT

    ii. Metal is conductor because of mobility of valence electronsiii. Good heat conductor because mobile electrons can carry additional

    kinetic energy across the metal

    c. Molecular orbital theory of metalsi. Molecular orbitals- form between two atoms when atomic orbitals on

    the atoms overlapii. Some cases, atomic orbitals on three or more atoms overlap to form

    molecular orbitals that encompass all atoms (these molecular orbitalsare delocalized)

    iii. Number of molecular orbitals that form by overlap of atomic orbitals= number of atomic orbitals

    iv. Metal: outer orbitals of an enormous number of metal atoms overlapto form an enormous number of molecular orbitals that aredelocalized over the metals a large number of energy levels arecrowded together into bands band theory

    v. Band theory (sodium): [Ne]3s11. When two Na atoms approach each other their 3s orbitals

    overlap to form two molecular orbitals when N atomsapproach each other there is N molecular orbitals overlappingand encompassing the whole crystal

    2. At each stage the number of energy levels grows eventuallythe energy levels effectively merge into a band of continuousenergies (called the 3s band for sodium metal)

    3. 3s band formed of N atoms will have N orbitals that hold a maxof 2N electrons

    4. Sodium has one valence electron, N atoms will supply N 3selectrons and will half-fill the 3s band

    5. Conductivity because electrons become free to movethroughout a crystal when they are excted to unoccupiedorbitals of a band requires very little energy in sodium(because of half filled 3s band, unoccupied orbitals lie just above the occupied orbitals of highest energy)

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    6. When voltage applied, electrons are excited into theunoccupied orbitals and move towards positive pole of voltagesource CURRENT

    vi. Band theory (magnesium): [Ne]3s21. Forms a 3s band like sodium

    2. Both sodium and magnesium have an unoccupied 3p bandformed from the unoccupied 3p orbitals3. Presence of this band in sodium = no effect 4. Magnesium- orbitals of individual atoms interact forming the

    metals, the energy levels spread so that bottom of 3p bandmerges with top of 3s band

    5. When electrons reach energy where two bands have merged,electrons begin to fill orbitals in both bands 3s and 3p bandsof magnesium metal are only partially filled by the time youhave accounted for all 2N valence electrons

    6. Voltage highest energy electrons are easily excited to

    unoccupied orbitals, CONDUCTORvii. IN GENERAL: solid that has a partially filled band will be an electrical

    conductor; solid that has only completely filled bands (without anearby unfilled band) will be a nonconductor (insulator)

    4. Group IA: The Alkali Metalsa. Except hydrogen, are all soft silvery metalsb. Most reactive of all metals (react readily with air and water)c. Ns1: usually form +1 ions (most compounds of these ions are soluble in

    water)d. Never occur as free metals in naturee. Lithium

    i. Chemically reactiveii. Relatively soft, hardest of all Group 1A

    iii. Many of ionic compounds are less soluble than similar compounds of other alkali metals

    iv. Metallurgy: source = the ore spodumenev. Reactions:

    1. Reacts readily with water and with moisture in air to producelithium hydroxide and hydrogen gas (not as vigorous asmixture of sodium and water)

    2. Burns in air to produce lithium oxide3. Reacts with nitrogen gas to form lithium nitride

    vi. Lithium compounds:1. Lithium carbonate: Li2CO3 primary source of all other lithium

    compounds2. Lithium hydroxide: LiOH strong base produced by reaction of

    lithium carbonate with calcium hydroxide in a precipitationreaction

    a. Li2CO3 + Ca(OH) 2 2LiOH + CaCO 3 f. Sodium and Potassium

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    i. Sodium metal = more reactive than potassium metalii. As you move down column of alkali metals, metals become more

    chemically reactive (forming the +1 ions)1. Partly because of decrease in ionization energy2. Atomic size increases going from top to bottom, valence

    electrons held less stronglyiii. Group 1A metals also increase in softness going from top to bottom1. Result of increasing atomic size, which results in decreasing

    strength of bonding by valence s shell electronsiv. Sodium metallurgy: obtained by electrolysis of molten sodium

    chloride1. Strong reducing agent

    v. Reactions of Sodium Metal:1. Reacts with water (greater vigor than lithium)2. Burns in air, producing some sodium oxide but mainly sodium

    peroxide

    a. Peroxide ion = oxidizing agent vi. Sodium compounds:

    1. NaOH one of top ten industrial chemicals; strong base2. Sodium carbonate= important (Na 2CO3) consumed with sand

    and lime in making glassvii. Potassium and potassium compounds

    1. Principal source of potassium and potassium compounds is KCl2. Almost all of potassium metal produced is used in prep of

    potassium superoxide KO 2 for self-contained breathingapparatus

    5. Group IIA: The Alkaline Earth Metalsa. Less reactive and harder than Alkali metalsb. Occur in nature as silicate rocks/ carbonates/ sulfatesc. Magnesium

    i. Metallurgy:1. Important as a structural metal

    a. Low density and relative strength of alloys2. Pure magnesium is relatively reactive (alloys add strength and

    corrosive resistance)3. Magnesium metal used as a reducing agent in the manufacture

    of titanium and zirconium from their tetrachloridesii. Reactions of magnesium metal

    1. Burns in air vigorously to form the oxide2. Burns in carbon dioxide, producing magnesium oxide and soot,

    or carbon3. Pure metal reacts slowly with water but reacts readily with

    steamiii. Magnesium compounds

    1. When magnesite, a magnesium carbonate mineral, is heatedabove 350, it decomposes to the oxide, a white solid

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    a. Careful heating at low temp results in a powdery formof the oxide that is relatively reactive reacts slowlywith water to produce magnesium hydroxide but reactsreadily with acids to yield the corresponding salts

    d. Calcium

    i. Metallurgy: obtained by reduction of calcium oxide with aluminum1. Mainly used in some alloysii. Reactions of calcium metal

    1. Soft reactive metal2. Reacts with water like the alkali metals to produce the metal

    hydroxide and hydrogen, the reactions are much less vigorous3. Calcium burns in air to produce the oxide CaO, and with

    chlorine to produce calcium chloride calcium reacts directlywith hydrogen to give the hydride

    iii. Calcium compounds1. Limestone, dolomite, anhydrite, gypsum

    2. Calcium oxide= in top ten industrial chemicals6. Group IIIA and Group IVA Metals

    a. Group IIIA- increasing metallic character in going down the columni. Boron- metalloid- chemistry is typical of nonmetal

    1. B(OH) 3 is actually acidicii. Rest of the elements are metals, but their hydroxides change from

    amphoteric for aluminum and gallium to basic for indium andthallium

    iii. Bonding to Boron is covalent, atom shares its 2s2sp1 valenceelectrons to give +3 oxidation state

    iv. Aluminum has many covalent compounds, but also has definitely ionicones (AlF 3)

    1. Al+3 is present in aq solutions of aluminum saltsv. Gallium, indium, and thallium also have ionic compounds and give

    cations in aq solutionsvi. Aluminum: only +3 important

    vii. Gallium and indium: +1 compoundsviii. Thallium: +3 and +1 are important

    ix. Group IVA show trend of greater metallic character in going downcolumn

    1. Last two in column, tin and lead, well known metalsa. Exist in +2 and +4 statesb. Tin: both are commonc. Lead: mostly +2

    b. Aluminum i. Metallurgy: soft and chemically reactive; very good conductor of

    electricityii. Reactions of aluminum

    1. Much less reactive than alkali and alkaline earth2. Does not react at an appreciable rate with water at room temp

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    3. Reacts readily with oxygen, but oxide coating that forms makesit corrosion resistant

    iii. Aluminum compounds1. Most important= aluminum oxide (alumina) Al 2O3

    a. Most used to make aluminum metals, some used as a

    carrier (support) for many heterogeneous catalystsrequired in chemical processesc. Tin and Lead

    i. Tin- two different forms (allotropes) one is metal and the other isnonmetal

    1. Nonmetal called gray tin (gray powder)2. Metallic- white tin (at low temp white tin transitions to gray

    tin)ii. Reactions of Tin and Lead

    1. Much less reactive than other metals of group IA IIA and IIIA2. Aluminum reacts vigorously with dilute hydrochloric and

    sulfuric acids, tin reacts only slowly with these acids3. Tin reacts more rapidly with the concentrated acids4. Lead metal reacts with these acids, but the products are

    insoluble and adhere to the metal so the reaction soon stopsiii. Tin and Lead compounds

    1. Tin(II) Chloride, SnCl 2, used as a reducing agent 2. Lead exists in compounds in the +2 and +4 states but Lead (II)

    more common

    Chemistry of Non-metals

    Hydrogen1. Properties and preparation of hydrogen

    a. Easily produced on a small scale:i. 2HCl(g) + Zn(s) ZnCl2(aq) + H 2(g)

    b. Colorless, odorless gas, less dense than airc. Not much in common with Group IA elements

    i. Much less likely to form a cationii. Forms covalent compounds with nonmetallic elements

    1. CH4, H2S, PH3 d. Three isotopes

    i. Protium, deuterium, and tritium1. Deuterium has boiling point of 101.42 degrees Celsius2. Tritium is radioactive

    ii. Often used as a marker in experimentse. Hydrogen is generally prepared industrially by steam-reforming process

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    i. Steam with hydrocarbons from natural gas or petroleum react at hightemperature and pressure in the presence of a catalyst to form carbonmonoxide and hydrogen

    ii. Water-gas reaction steam is passed over red-hot coal, producingcarbon monoxide and hydrogen gas

    iii. To remove carbon monoxide, it is reacted with steam in the presenceof a catalyst to give CO 2 and more hydrogen gasiv. The CO2 is then removed by passing the mixture of gases through a

    basic aqueous solution2. Hydrogen reactions and compounds

    a. Mostly used to prepare ammoniab. Also used in petrochemical industryc. Food industry

    i. Hydrogenation converts liquid fats to solid fats by saturating carbon-carbon double bonds with H 2 gas

    d. Synthesis gas reaction

    i. Cobalt-catalyzed reaction with carbon monoxide to produce methanole. Reduction of metal oxides to extract pure metalsf. Binary hydride

    i. A compound that contains hydrogen and one other element ii. Ionic hydrides

    1. Contain hydride ion (H -)2. Can be directly forms via reaction of an alkali metal or large

    Group IIA metals (Ca, Sr, and Ba) with H 2 gas near 400 degreesCelsius

    a. 2Li(s) + H 2(g) 2LiH(s)b. Ba(s) + H 2(g) BaH2(g)

    3. White crystalline compounds4. H has oxidation state of -15. Can undergo oxidation-reduction reaction with water to

    produce hydrogen and a basic solutiona. LiH(s) + H 2O(l) H2(g) + LiOH(aq)

    iii. Covalent hydrides1. Molecular compounds in which hydrogen is covalently bonded

    to another element a. NH3, H2O, H2O2

    2. Some formed by direct reaction of the elements3. If the nonmetal reacting with hydrogen is reactive, the reaction

    will readily occur without high temp. or a catalayst a. F2(g) + H 2 (g) 2HF(g)

    4. Other reactions will require temp. or catalyst a. 2H 2(g) + O 2(g) 2H2O(g) [exothermic, rapid]

    5. Clean fuel source, most efficient heat per gram of any fueliv. Metallic hydrides

    1. Compounds containing transition metal and hydrogen2. MHx, where x is often not an integer

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    3. Non-stoichiometric amounts of hydrogen enter a crystallinestructure under different conditions, e.g. pressure

    a. MH0.4 , MH0.7

    Group IVA the Carbon Family

    3. Carbona. Catenation the covalent bonding of two or more atoms of the same element to one another

    i. Carbon is by far the most prevalent b. Allotropes of carbon

    i. Diamond1. Each carbon tetrahedrally (sp 3) bonded to four other carbon

    atoms2. Among hardest substances known

    ii. Graphite1. Black substance

    2. Layer structurea. Each layer consists of carbon atoms bonded to three

    other carbon atoms to give hexagonal pattern of carbonatoms arranged in a plan

    b. Layers easily slide over one another, making it soft 3. Bonding involved sp 2 hybridization of the carbon atoms with

    delocalized pi electrons4. Good conductor, because of delocalized bonding within layers

    iii. Buckminsterfullerene1. Stable soccer ball form

    c. Carbon black i. Composed of extremely small crystals of carbon with imperfect

    graphite structureii. Used in tires, printer ink

    d. Oxides of carboni. Carbon and organic compounds burn in an excess of O 2 to give CO 2

    ii. Equilibrium among carbon, carbon dioxide, and carbon monoxidefavors CO above 700 degrees Celsius

    1. CO2(g) + C(s) 2CO(g)2. CO is almost always a product of combustion

    iii. CO is produced industrially from CH 4 with steam or partial oxidation1. CH4(g) + H 2O(g) CO(g) + 3H 2(g)

    iv. Methanol production requires CO1. CO(g) + 2H 2(g) CH3OH

    v. CO2 is colorless, odorlessvi. CO2 is produced whenever carbon or organic materials are burned

    e. Carbonatesi. CO2 dissolves in water to form aqueous solution of carbonic acid

    1. CO2(g) + H 2O(l) H2CO3(aq)

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    ii. Carbonic acid is diprotic (two acidic H atoms), and dissociates to formcarbonate ion

    iii. Carbonate ion then dissociates to form hydrogen carbonate andcarbonate ions

    iv. CO2 reacts very readily to form salts

    1. CO2(g)+ Ca(OH) 2(aq) CaCO3(s) + H 2O(l)4. Silicona. Basic material in semiconductor devices in most electronicsb. Silicates compounds of silicon and oxygen with one or more metallic

    compoundsc. Elemental silicon is obtained by reducing quartz san (SiO 2) with coke (C) at

    3000 degrees Celsiusi. SiO2(l) + 2C(s) Si(l) + 2CO(g)

    d. Diamond-like structuree. Must be exceptionally pure for industrial use

    i. Si(s) + 2Cl 2(g) SiCl4(l)

    ii. SiCl4(g) + 2H 2(g) SiCl4(l)1. Pure silicon crystallizes on surface of pure silicon rod

    f. Silica (silicon dioxide)i. SiO2

    ii. A covalent network solid in which each silicon atom is covalentlybonded in tetrahedral directions to four oxygen atom, and eachoxygen atom is in turn bonded to another silicon atom

    iii. Quartz crystals1. Piezoelectric effect

    a. In a piezoelectric crystal, like quartz, compression of thecrystal in a particular direction causes an electricvoltage to develop across it

    b. This effect allows quartz crystals to be used in soundsystems, as they are sensitive to sound vibrations andcan convert them to electrical currents

    g. Silicatesi. A compound of silicon and oxygen with one or more metals that may

    be formally regarded as a derivative of silicic acid, H 4SiO4 or Si(OH) 4 ii. Not isolated, but silicate ions are present in liquid silica

    1. Via condensation reactions , two molecules of silicate arechemically joined by the elimination of H 2O

    2. All silicate minerals have long chain structureh. Silicones

    i. A polymer containing chains of silicon-oxygen bonds, withhydrocarbon groups such as CH 3 attached to silicon atoms

    ii. Preparation:1. Reaction of silicon with methyl chloride at high temp with

    copper catalyst a. Si(s) + 2CH 3Cl(g) (CH3)SiCl2

    2. The product is reacted with water

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    a. (CH3) 2SiCl2(l) + 2H 2O(l) (CH3) 2Si(OH) 2 + 2HCl(g)3. Then the product undergoes a condensation reaction to form a

    silicone oil

    Group VA Nitrogen and Phosphorous Family

    5. Nitrogena. N2 makes up 78.1% of the atmospherei. Relatively unreactive

    b. Exists in important compounds with oxidation states between -3 and +5c. Properties and uses

    i. Stable triple bondii. Some very reactive metals do react with N 2

    1. Magnesium metal burns in air, forming nitride as well as oxidea. 3Mg(s) + N 2(g) Mg3N2(s)

    2. N3- is very strong base, reacting with water to produceammonia

    a. N3-(aq) + 3H 2O(l) NH3(g) + 3OH -(aq)iii. Good blanketing gas, which is used as the atmosphere for protecting

    a material from oxygen in the air during storaged. Nitrogen compounds

    i. Ammonia (NH 3)1. Colorless gas2. Prepared from N 2 and H 2 by Haber process 3. Can be prepared in the lab with ammonium salt and a strong

    basea. NH4Cl(aq) + NaOH(aq) NH3(g) + H 2O(l) + NaCl(aq)

    4. Liquid used as fertilizerii. Nitrous oxide (N 2O)

    1. Prepared by careful heating of molten ammonium nitratea. NH4NO3(s) N2O(g) + 2H 2O(g)

    iii. Nitrogen monoxide (NO) aka nitric oxide 1. In preparation, ammonia is oxidized in the presence of

    platinum catalyst a. 4NH 3(g) + 5O 2(g) 4NO(g) + 6H 2O(g)

    iv. Nitric acid (HNO 3)1. Produced by Ostwald process

    a. Ammonia is burned in the presence of platinum catalyst to give NO

    b. NO reacts with oxygen to form NO 2 c. NO2 is dissolved in water, where it reacts to form nitric

    acid and NO2. Strong oxidizing agent

    a. 3Cu(s) + 8H 3O+ (aq) + 2NO 3- 3Cu 2+ (aq) + 2NO(g) +12H 2O(l)

    v. When certain nitrates are heated, they decompose to the nitrites1. 2NaNO 3(s) 2NaNO 2(s) + O 2(g)

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    2. The corresponding acid, nitrous acid (2HNO 2), is unstable6. Phosphorous

    a. Exists in important compounds with oxidation states of +3 to +5b. Allotropes of phosphorous

    i. White phosphorous (P 4)

    1. Waxy, white solid2. Highly unstable and reactive3. Molecular solid

    a. A solid held together by van der Waals forcesb. Low melting point

    4. Forms tetrahedral pyramid with high-energy 60-degree bondangles

    a. Seeks to replace bonds with more stable ones5. 2Ca 3(PO 4) 2(s) + 6SiO 2(s) + 10C(s) 6CaSiO3(l) + 10CO(g) + P 4(g)

    a. Gases are cooled by water to condense phosphorousvapor to the liquid

    ii. Red phosphorous1. Made by heating white phosphorous to 400 degrees C2. Relatively safe, amorphous

    c. Phosphorous oxidesi. Tetraphosphorus hexoxide (P 4O6)

    1. Tetrahedron of phosphorous atoms with oxygen atomsbetween each pair of P atoms

    2. Low-melting solid (23 degrees C)ii. Tetraphosphorus decoxide (P 4O10 )

    1. Structure similar to P 4O6, except each P has an additional Obonded to it

    2. Preparation by burning white phosphorous in aira. P4(s) + 5O 2(g) P4O10

    d. Phosphorous oxoacidsi. Phosphoric acid

    1. H3PO4 2. Salts: NaH 2PO4, Na2HPO4, Na3PO4 3. Obtained from tetraphosphorus decoxide or from rock

    a. Ca3(PO 4) 2 + 3H 2SO4(aq) 3CaSO4(s) + 2H 3PO4(aq)4. Undergoes condensation reactions to form other phosphoric

    acids (diphosphoric acid, triphosphoric acid, et cetera)a. Hn+2 PnO3n+1

    ii. Metaphosphoric acids1. Acids with the general formula (HPO 3) n

    a. When n is very large, the acid is calledpolymetaphosphoric acid

    iii. Cause eutrophication of aquatic ecosystems

    Group VIA Oxygen and sulfur family1. Oxygen

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    a. Exists mainly in -2 oxidation statei. Other Group VIA elements exist mostly in +4 to +6 state

    b. Very electronegativec. Bonding involves on the s and p orbitals

    i. For other elements in Group VIA, d orbitals are a factor in bonding

    d. Most oxygen on earth is present as oxide or oxoanion mineralsi. Only 20% of Earths oxygen exists as O 2 in the aire. Properties and preparation of oxygen

    i. Common form is O 2 ii. Very low-melting and low-boiling

    iii. Can be prepared in small quantities by decomposing certain oxygen-containing compounds

    1. Heating mercury oxidea. 2HgO(s) 2Hg(l) + O 2(g)

    2. Heating KClO 3 with pure manganese(IV) oxide catalyst a. 2KClO3(s) 2KCl(s) + 3O 2(g)

    f. Reactions of oxygeni. Molecular oxygen is very reactive

    ii. Oxide binary compound with oxygen in the -2 oxidation state1. Most metals react readily with oxygen to form oxides,

    especially metals in a form with large surface area2. 2Mg(s) + O 2(g) 2MgO(s)3. 3Fe(s) + 2O 2 Fe3O4(s)

    iii. Metal oxides are basic oxides1. Metals in a high oxidation state may be acids

    a. Cr2O3 is a baseb. CrO3 is an acid

    iv. Alkali metal oxides1. When burned, yield peroxides

    a. A compound with oxygen in the -1 oxidation stateb. 2Na(s) + O 2(g) Na2O2(s)

    2. When burned, yield superoxides a. Binary compound with oxygen in the -1/2 oxidation

    stateb. K(s) + O 2 KO2(s)

    v. Nonmetals react with oxygen to form covalent oxides, which aremostly all acidic

    1. C burns in excess O 2 to form CO 2 2. S8(s) + 8O 2(g) 8SO2(g)

    vi. Compounds with at least one element in a reduced state are oxidizedby oxygen, giving compounds that would be expected to form whenthe individual elements are burned in oxygen

    1. C8H18 burns in oxygen to give CO 2 and watera. 2C8H18 (l) + 25O 2(g) 16CO 2(g) + 18H 2O(g)

    2. 2H2S(g) + 3O 2(g) 2H2O(g) + 2SO 2(g)2. Sulfur

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    a. Free sulfur formed by volcanic gas reactionsi. 16H 2S(g) + 8SO 2(g) 16H 2O(l) + 3S 8(s)

    b. Allotropes of sulfuri. Rhombic sulfur most stable allotrope of sulfur under normal

    conditions

    1. Yellow, crystalline lattice of crown-shaped S 8 moleculesii. Monoclinic sulfur same structure as rhombic sulfur, except for theway the molecules are packed to form crystals

    iii. Unstable below 96 degrees C, reverts back to rhombic sulfurc. Production of sulfur

    i. Mined from deep underground deposits by Frash processii. Hydrogen sulfide, H 2S, recovered from natural gas and petroleum

    1. Also a source of free sulfur2. Free sulfur obtained from H 2S by Claus process

    a. 8H 2S(g) + 4O 2(g) S8(s) + 8H 2O(g)b. 2H 2S(g) + 3O 2(g) 2SO2(g) + 2H 2O(g)

    c. 16H 2S(g) + 8SO 2(g) 3S8(s) + 16H 2O(g)d. Sulfur oxides and oxoacids

    i. SO2 1. Suffocating odor2. Present in polluted air3. Very soluble in water, producing acidic solutions

    a. H2O(l) + SO 2(g) H2SO3(aq)b. H2SO3(aq) + H 2O(l) H3O+ (aq) + HSO 3 c. HSO3(aq) + H 2O(l) H3O+ (aq) + SO 32(aq)

    4. With appropriate amount of base, the corresponding hydrogensulfite salt or sulfite salt is obtained

    a. Na2CO3(aq) + 2SO 2(aq) + H 2O(l) 2NaHSO 3(aq) + CO 2(g)b. Na2CO3(aq) + SO 2(aq) Na2SO3(aq) + CO 2(g)

    5. When treated with acids, sulfites and hydrogen sulfitesdecompose to produce SO 2

    a. NaHSO3 + HCl(aq) NaCl(aq) + H 2O(l) + SO 2(g)6. Burning S 8 produces SO 2

    ii. SO3 1. Exists in liquid form at room temp in equilibrium with S 3O9 2. Vapor-phase is is SO 3, planar triangular geometry3. SO2 reacts slowly with oxygen in air to form SO 3

    a. Reaction is much faster with catalyst, e.g. platinumb. 2SO2(g) + O 2(g) 2SO3(g)

    4. SO3 reacts vigorously and exothermically with water toproduce sulfuric acid

    a. SO3(g) + H 2O(l) H2SO4(aq)iii. Contact process

    1. Industrial method for synthesis of sulfuric acid

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    2. Reaction of sulfur dioxide with oxygen to form sulfur trioxideusing vanadium oxide catalyst, followed by reaction of sulfurtrioxide with water

    iv. SO3 in the atmosphere dissolves in rain, producing H 2SO3, and fallingas acid rain

    v. Hot, concentrated sulfuric acid can oxidize metals not normaldissolved by acids, like copper

    Group VIIA the Halogense. Fluorides differ somewhat from the rest of the group

    i. Not soluble in water, whereas CaCl, CaBr, and Iodide are very solublein water

    ii. All, including fluorides, form stable compounds in which the halogenis in a -1 oxidation state

    1. In fluorine compounds, this is the only oxidation state2. In Cl, Br, and I, compounds can form in which the halogen is in

    a positive oxidation state (+1, +3, +5, +7)3. Higher positive oxidation states are due to d orbitals in

    bonding3. Chlorine

    a. Discovered in 1774 by heating HCl with manganese oxidei. 4HCl(aq) + MnO 2(s) MnCl2(aq) + Cl 2(g) + 2H 2O(l)

    b. Most commercially important halogenc. Very reactive oxidizing agent, similar to oxygen

    i. Oxidizing power decreases among halogens down the groupii. Chlorine is stronger oxidizing agent than bromine or iodide

    1. Cl2(g) + 2KBr(aq) 2KCl(aq) + Br 2(aq)2. Cl2(g) + 2KI(aq) 2KCl(aq) + I 2(aq)

    d. Chlorine reacts with water by being both oxidized and reducedi. Cl2(g) + H 2O(l) HClO(aq) + HCl(aq)

    e. Hydrogen chloridei. Very soluble in water

    1. HCl(g) + H 2O(l) H3O+ (aq) + Cl (aq)ii. HCl can be produced by heating NaCl with concentrated sulfuric acid

    1. NaCl(s) + H 2SO4(l) NaHSO4(s) + HCl(g)iii. Heated more strongly, sodium hydrogen sulfate reacts with sodium

    chloride to produce additional hydrogen chlorida1. NaCl(s) + NaHSO 4(s) Na2SO4(s) + HCl(g)2. Hydrogen bromide and hydrogen iodide can prepared by

    analogous reactions replacement reaction with their saltsa. Phosphoric acid used instead of HCl because sulfuric

    acid will oxidize bromide and iodide ionsb. 2Br (aq) + HCl(aq) Br2(g)

    3. HCl prepared as byproduct of chlorinate hydrocarbona. CH4(g) + Cl 2(g) CH3Cl(g) + HCl(g)

    f. Oxoacids of chlorine

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    i. More acidic with more oxygen atoms bonded to halogen atom1. (weakest) HClO, HClO 2, HClO3, HClO4 (strongest)

    ii. Hypochlorous acid1. Cl2(g) + 2OH (aq) Cl(aq) + ClO (aq) + H 2O(l)2. Hypochlorite ion is unstable, and decomposes

    a. 3ClO

    (aq) ClO3

    (aq) + 2Cl

    (aq)3. The reaction is fast with heat and a basea. 3Cl2(g) + 6NaOH(aq) NaClO3(aq) + 5NaCl(aq) + 3H 2O(l)

    4. Production of percholate and potassium percholatea. ClO3(aq) + 3H 2O(l) ClO4(aq) + 2H 3O+ (aq) + 2e b. KClO4(s) + H 2SO4(l) KHSO4(s) + HClO 4(l)

    Group VIIIA the Noble Gases

    4. Helium and the other Noble Gasesa. All except helium and radon are obtained by distillation of liquid air

    b. Helium has lowest boiling point (-268.9 degrees Celsius) of any substancec. All noble gases are used in gas discharge tubesd. Argon is mixed with nitrogen to fill incandescent lightbulbse. Neon used in signs because of highly visible red-orange emissionf. Xenon reacts directly with fluorine at 400 degrees Celsius to give

    tetrafluoridei. Xe(g) + 2F 2(g) XeF4(s)

    ii. Reaction at room temperature when exposed to sunlight