Periodic Table It is a systematic catalog of the elements. Elements are arranged in order of atomic number.

Periodicity When one looks at the chemical properties of elements, one notices a repeating pattern of reactivities.

Periodic Table The rows on the periodic chart are periods. Columns are groups. Elements in the same group have similar chemical properties. Chemists who knew nothing about electrons developed the table- meant to correlate behaviors of elements and to help us remember many facts

Groups These five groups are known by their names.

Periodic Table Nonmetals are on the right side of the periodic table (with the exception of H). Gases – Group 18 and F, O , N , H, Cl Liquid – Br Rest are solids

Periodic Table Metalloids border the stair-step line (with the exception of Al, Po, and At).

Periodic Table Metals are on the left side of the chart. Properties – luster and high electrical and heat conductivity malleable All except Hg are solids at room temperature

Development of Periodic Table Elements in the same group generally have similar chemical properties. Physical properties are not necessarily similar, however.

Development of Periodic Table Dmitri Mendeleev and Lothar Meyer independently came to the same conclusion about how elements should be grouped.

Development of Periodic Table Mendeleev, for instance, predicted the discovery of germanium (which he called eka-silicon meaning under silicon) as an element with an atomic weight between that of zinc and arsenic, but with chemical properties similar to those of silicon. He left spaces in the table which allowed him to boldly predicted existence and properties

Development of Periodic Table 1913 – Henry Mosley developed concept of atomic number Bombarded different elements with high-energy electrons, found each element produced X-rays of a unique frequency and that the frequency generally increased with the atomic mass He arranged the X-ray frequencies in order by assigning them a whole number and called it the atomic number He correctly identifies the atomic number as the number of protons in a nucleus He left spaces in the table which allowed him to boldly predicted existence and properties

Development of Periodic Table Concept of atomic number clarified some problems in the periodic table Ex) atomic weight of Ar (18) is greater than that of K (19) yet the chemical properties of Ar are much more like those of Ne and Kr than those of Na and Rb When elements are arranged by atomic number, rather than increasing atomic weight, Ar and K appear in the correct places This made is possible to discover previously unknown elements He left spaces in the table which allowed him to boldly predicted existence and properties

Periodic Trends In this chapter, we will rationalize observed trends in Sizes of atoms and ions. Ionization energy. Electron affinity.

Effective Nuclear Charge Coulombs Law states that the strength of the interaction between two electrical charges depends on the magnitude of the charges and the distances between them .

Effective Nuclear Charge The attractive force between an electron and the nucleus depends on the nuclear charge and on the average distance between the nucleus and the electron The force increases as the nuclear charge increases and decreases as the electron moves further away

Effective Nuclear Charge In a many-electron atom, electrons are both attracted to the nucleus and repelled by other electrons. The nuclear charge that an electron experiences depends on both factors. In general so many electron-electron repulsions that we cannot analyze the situation exactly

Effective Nuclear Charge Can treat each electron as though it were moving in the net electric field created by the nucleus and the electron density of the other electrons We view this net electric field as if it results from a single positive charge in the nuclues

Effective Nuclear Charge In many electron atoms the inner electrons partially screen outer electrons from the attraction of the nucleus

Effective Nuclear Charge The effective nuclear charge, Zeff, is found this way: Zeff = Z − S where Z is the atomic number and S (positive number) is a screening constant, usually close to the number of inner electrons. The effective nuclear charge is smaller than the actual nuclear charge because the effective nuclear charge includes the effect of the other electrons Screening constant represents the portion of the nuclear charge that is screened from the valence electrons by other electrons in the atom Electrons in the same valence shell do not screen one another very effectively, but they do affect the value of S slightly

Effective Nuclear Charge Na nuclear charge is Z= 11+ there are 10 core electrons 1s22s22p6 Expect S = 10 and the 3s electron to experience an effective nuclear charge of Zeff = 11-10 = 1+ Situation is more complicated – 3s electron has small probability of being closer to the nucleus in the region occupied by the core electrons – thus there is a possibility that this electron experiences an attraction greater than our simple S = 10 model suggests This greater attraction turns out to increase the value Zeff for the 3s electron in Na from our expected Zeff = 1+ to Zeff = 2.5+ The fact that the 3s electron spends some small amount of time close to the nucleus changes the value of S from 10 – 8.5

Effective Nuclear Charge For many-electron atom the energies of orbitals with the same n value increases with increasing l value In C 1s22s22p6, the energy of the 2p orbital (l=1) is higher than that of 2s (l=0) even though both orbitals are in the n=2 shell Difference in energies is due to the radial probability functions for the orbitals The greater attraction between 2s electron and the nucleus leads to a lower energy for the 2s orbital than for the 2 p orbital

Effective Nuclear Charge Trends in valence electron Zeff values: The effective nuclear charge increases from left to right across any period of the periodic table Number of core electrons stays the same across the period, protons increase Down a column the effective nuclear charge experienced by the valence electrons changes far less than it does across a period

Effective Nuclear Charge n is larger than the value of n for the electron of interest contributes 0 n is equal to the value of n for the electron of interest contribute 0.35 n is 1 less than the n value for the electron of interest contribute .85 n is even smaller contribute 1.00 Ex) F 1s22s22p5

Effective Nuclear Charge – Slater Ex) F 1s22s22p5 Electron of interest n = 2 n = 2 (7 electrons will look at the effect of 6 on 1 ) S = (6*.35) + (2*.85) = 3.8 Zeff = Z- S = 9 – 3.8 = 5.2+

What Is the Size of an Atom? The shortest distance separating the two nuclei during collisions is twice the radii of the atoms – non bonding atomic radius or van der Waals radius

What Is the Size of an Atom? The bonding atomic radius is defined as one-half of the distance between covalently bonded nuclei. – shorter than the nonbonding atomic radius The attractive force between two adjacent atoms in the molecule leading to a chemical bond causes this

What Is the Size of an Atom? Example: I2 molecule, the distance separating the nuclei is observed to be 2.66Å, which means the bonding atomic radius of an iodine atom is (2.66Å)/2 = 1.33Å Helium and neon, the bonding atomic radius must be estimated because there are no known compounds of these elements

What Is the Size of an Atom? Knowing atomic radii allows us to estimate bond lengths Ex) Cl-Cl bond length in Cl2= 1.99Å so bonding atomic radius is 0.99Ǻ In CCl4 bond length 1.77Ǻ bonding atomic radius Close to the sum (0.77+0.99) 1 Ǻ = 10-10m

Sizes of Atoms Bonding atomic radius tends to… …decrease from left to right across a row (due to increasing Zeff). …increase from top to bottom of a column (due to increasing value of n).

Sizes of Ions Ionic size depends upon: The nuclear charge. The number of electrons. The orbitals in which electrons reside.

Sizes of Ions Cations are smaller than their parent atoms. The outermost electron is removed and repulsions between electrons are reduced. MELPS Helps!!

Sizes of Ions Anions are larger than their parent atoms. Electrons are added and repulsions between electrons are increased.

Sizes of Ions Ions carrying the same charge ionic radius will increase in size as you go down a column. This is due to increasing value of n.

Sizes of Ions In an isoelectronic series, ions have the same number of electrons. Ionic size decreases with an increasing nuclear charge.

Ionization Energy The ionization energy is the amount of energy required to remove an electron from the ground state of a gaseous atom or ion. The first ionization energy is that energy required to remove first electron. The second ionization energy is that energy required to remove second electron, etc.

Ionization Energy It requires more energy to remove each successive electron. When all valence electrons have been removed, the ionization energy takes a quantum leap. Supports the idea that only the outermost electrons are involved in the sharing and transfer of electrons that give rise to chemical bonding and reactions

Trends in First Ionization Energies As one goes down a column, less energy is required to remove the first electron. For atoms in the same group, Zeff is essentially the same, but the valence electrons are farther from the nucleus.

Trends in First Ionization Energies Generally, as one goes across a row, it gets harder to remove an electron. As you go from left to right, Zeff increases.

Trends in First Ionization Energies However, there are two apparent discontinuities in this trend.

Trends in First Ionization Energies The first occurs between Groups IIA and IIIA. In this case the electron is removed from a p-orbital rather than an s-orbital. The electron removed is farther from nucleus. There is also a small amount of repulsion by the s electrons.

Trends in First Ionization Energies The second occurs between Groups VA and VIA. The electron removed comes from doubly occupied orbital. Repulsion from the other electron in the orbital aids in its removal.

Electron Configurations of Ions When electrons are removed from an atom to form a cation, they are always removed first from the occupied orbitals having the largest n value Electrons added to form an anion are added to the empty or partially filled orbital having the lowest value of n

Electron Affinity Electron affinity is the energy change accompanying the addition of an electron to a gaseous atom: Measures the attraction or affinity of the atom fro the added electron – negative sign would indicate that energy is released during the process Cl + e−  Cl−

Trends in Electron Affinity In general, electron affinity becomes more exothermic as you go from left to right across a row.

Trends in Electron Affinity There are again, however, two discontinuities in this trend.

Trends in Electron Affinity The first occurs between Groups IA and IIA. The added electron must go in a p-orbital, not an s-orbital. The electron is farther from nucleus and feels repulsion from the s-electrons.

Trends in Electron Affinity The second occurs between Groups IVA and VA. Group VA has no empty orbitals. The extra electron must go into an already occupied orbital, creating repulsion.

Properties of Metal, Nonmetals, and Metalloids © 2009, Prentice-Hall, Inc.

Metals versus Nonmetals Differences between metals and nonmetals tend to revolve around these properties.

Metals versus Nonmetals Metals tend to form cations. Nonmetals tend to form anions. MELPS

Metals They tend to be lustrous, malleable, ductile, and good conductors of heat and electricity.

Metals Compounds formed between metals and nonmetals tend to be ionic. Metal oxides tend to be basic.

Nonmetals These are dull, brittle substances that are poor conductors of heat and electricity. They tend to gain electrons in reactions with metals to acquire a noble gas configuration.

Nonmetals Substances containing only nonmetals are molecular compounds. Most nonmetal oxides are acidic.

Metalloids These have some characteristics of metals and some of nonmetals. For instance, silicon looks shiny, but is brittle and fairly poor conductor.

Metals Compounds formed between metals and nonmetals tend to be ionic. Metal oxides tend to be basic.

Metallic Character Metallic character decreases from left to right across a period Metallic character increases from top to bottom in a group

Metal Ions Metals lose electrons and become positive ions (cations) with a smaller radius than the parent ion They are easily oxidized Have low ionization energies

Metal Ions Alkali metals have 1+ and alkaline earth metal 2+ Transition metals do not follow a pattern Compounds between metal and nonmetal are ionic

Alkali Metals Alkali metals are soft, metallic solids. The name comes from the Arabic word for ashes.

Alkali Metals They are found only in compounds in nature, not in their elemental forms. They have low densities and melting points. They also have low ionization energies.

Alkali Metals Their reactions with water are famously exothermic.

Alkali Metals Alkali metals (except Li) react with oxygen to form peroxides. K, Rb, and Cs also form superoxides: K + O2  KO2 They produce bright colors when placed in a flame. © 2009, Prentice-Hall, Inc.

Alkaline Earth Metals Alkaline earth metals have higher densities and melting points than alkali metals. Their ionization energies are low, but not as low as those of alkali metals.

Alkaline Earth Metals Beryllium does not react with water and magnesium reacts only with steam, but the others react readily with water. Reactivity tends to increase as you go down the group.

Alkaline Earth Metals The heavier alkaline earth ions give off characteristic colors when heated in a hot flame.

Nonmetals Substances containing only nonmetals are molecular compounds. Most nonmetal oxides are acidic.

Nonmetals These are dull, brittle substances that are poor conductors of heat and electricity. They tend to gain electrons in reactions with metals to acquire a noble gas configuration.

Group 6A Oxygen, sulfur, and selenium are nonmetals. Tellurium is a metalloid. The radioactive polonium is a metal.

HYDROGEN Located above be alkali metals because of its electron configuration Does not truly belong to any group - it is a nonmetal High ionization energy

Oxygen There are two allotropes of oxygen: There can be three anions: O3, ozone There can be three anions: O2−, oxide O22−, peroxide O21−, superoxide It tends to take electrons from other elements (oxidation).

Sulfur Sulfur is a weaker oxidizer than oxygen. The most stable allotrope is S8, a ringed molecule.

Group VIIA: Halogens The halogens are prototypical nonmetals. The name comes from the Greek words halos and gennao: “salt formers”.

Group VIIA: Halogens They have large, negative electron affinities. Therefore, they tend to oxidize other elements easily. They react directly with metals to form metal halides. Chlorine is added to water supplies to serve as a disinfectant

Group VIIIA: Noble Gases The noble gases have astronomical ionization energies. Their electron affinities are positive. Therefore, they are relatively unreactive. They are found as monatomic gases.

Group VIIIA: Noble Gases Xe forms three compounds: XeF2 XeF4 (at right) XeF6 Kr forms only one stable compound: KrF2 The unstable HArF was synthesized in 2000.

Metals Valence electrons are in a partially-filled band. © 2009, Prentice-Hall, Inc.

Metals metallic bonding results from the fact that the valence electrons are delocalized throughout the entire solid. Visualize this as an array of positive ions immersed in a “sea” of delocalized valence electrons Metals have luster, high thermal conductivity and high electrical conductivity Malleable and ductile

Structures of Metallic Solids Crystal structures of metals are simple enough we can generate the structure by placing a single atom on each lattice point Primitive cubic structure are rare – radioactive polonium Body-centered cubic metals – iron, chromium, sodium, and tungsten Face-centered cubic metals – aluminum, lead, copper, silver, and gold

CLOSE PACKING Shortage of valence electrons and the fact that they are collectively shared make it favorable for the atoms in a metal to pack together closely Atoms are spherical – understand structure of metals by looking at how spheres pack together Efficient way – pack one layer of equal sized spheres to surround each sphere by six neighbors For 3-D structure keeps stacking layer Second layer will sit in the depressions made by the first spheres

CLOSE PACKING Third layer – we have 2 choices First – third layer in the depressions that lie directly over the spheres in the first layer Fourth layer would lie directly over the spheres in the second layer (FIG 12-13) HEXAGONAL CLOSE PACKING (HCP)

CLOSE PACKING Second – third layer does not sit directly above the spheres in either of the first two layers, subsequent layers repeat this sequence giving an ABCABC pattern CUBIC CLOSE PACKING (CCP)

CLOSE PACKING In both each sphere has 12 equidistant nearest neighbors - 6 in the same layer 3 from the layer above and 3 from the layer below Each sphere has a coordinate number of 12 – the number of atoms immediately surrounding a given atom in a crystal structure

Alloys Alloys contain more than one element and have the characteristic properties of metals. Solid Solution alloys are homogeneous mixtures. Pure metals and alloys have different physical properties. An alloy of gold and copper is used in jewelry (the alloy is harder than the relatively soft pure 24 karat gold). 14 karat gold is an alloy containing 58% gold.

Metal Alloys-Solid Solutions Substance has mixture of element and metallic properties. 1. Substitutional Alloy: some metal atoms replaced by others of similar size. Electronegativities usually are similar. The atoms must have similar atomic radii. The elements must have similar bonding characteristics. brass = Cu/Zn

Metal Alloys (continued) 2. Interstitial Alloy: Interstices (holes) in closest packed metal structure are occupied by small atoms. Solute atoms occupy interstices “small holes” between solvent atoms. One element (usually a nonmetal) must have a significantly smaller radius than the other (in order to fit into the interstitial site). steel = iron + carbon 3. Both types: Alloy steels contain a mix of substitutional (Cr, Mo) and interstitial (Carbon) alloys.

Alloys vs. Pure Metal The alloy is much harder, stronger, and less ductile than the pure metal (increased bonding between nonmetal and metal). An example is steel (contains up to 3% carbon). mild steels (<0.2% carbon) useful for chains, nails, etc. medium steels (0.2-0.6% carbon) useful for girders, rails, etc. high-carbon steels (0.6-1.5% carbon) used in cutlery, tools, springs. Other elements may also be added to make alloy steels. Addition of V and Cr increases the strength of the steel and improves its resistance to stress and corrosion. The most important iron alloy is stainless steel. It contains C, Cr (from ferrochrome, FeCr2), and Ni. Heterogeneous alloys: The components are not dispersed uniformly (e.g., pearlite steel has two phases: almost pure Fe and cementite, Fe3C).

Substitutional Alloy Interstitial Alloy

Ordered intermetallic structures. Keywords: compounds with definite properties and composition attractive for high temperature Ni3Al used for component of jet aircraft because of the high temperatures and low density (a) The face-centered cubic unit cell of a cubic close-packed metal. (b) The ordered structure of Ni3Al with nickel atoms shown in gray and aluminum atoms in blue. (c) The ordered structure of TiAl with titanium atoms shown in red and aluminum atoms in blue.

Figure 23.17abcd Figure 23-17 Title: Ordered intermetallic structures. Caption: (a) The face-centered cubic unit cell of a cubic close-packed metal. (b) The ordered structure of Ni3Al with nickel atoms shown in gray and aluminum atoms in blue. (c) The ordered structure of TiAl with titanium atoms shown in red and aluminum atoms in blue. (d) The ordered structure of ZnCu with copper atoms shown in copper and zinc atoms in gray. Notes: Keywords:

Which two substances are most likely to form an interstitial alloy? Nickel and titanium Silver and tin Tin and lead Copper and zinc Tungsten and carbon

Which two substances are most likely to form an interstitial alloy? Nickel and titanium Silver and tin Tin and lead Copper and zinc Tungsten and carbon

Bonding Models for Metals Electron Sea Model: A regular array of metals in a “sea” of electrons. The electron-sea model is a qualitative interpretation of band theory (molecular-orbital model for metals). Band (Molecular Orbital) Model: Electrons assumed to travel around metal crystal in MOs formed from valence atomic orbitals of metal atoms. Conduction Bands: closely spaced empty molecular orbitals allow conductivity of heat and electricity.

Molecular-Orbital Model for Metals Delocalized bonding requires the atomic orbitals on one atom to interact with atomic orbitals on neighboring atoms. Example: Graphite electrons are delocalized over a whole plane, while benzene molecules have electrons delocalized over a ring. Recall that the number of molecular orbitals is equal to the number of atomic orbitals. Each orbital can hold two electrons. In metals there are a very large number of orbitals. As the number of orbitals increases, their energy spacing decreases and they band together. The available electrons do not completely fill the band of orbitals.

Molecular-Orbital Model for Metals Therefore, electrons can be promoted to unoccupied energy bands. Because the energy differences between orbitals are small the promotion of electrons requires little energy. As we move across the transition metal series, the antibonding band starts becoming filled. Therefore, the first half of the transition metal series has only bonding-bonding interactions and the second half has bonding–antibonding interactions. We expect the metals in the middle of the transition metal series (group 6B) to have the highest melting points. The energy gap between bands is called the band gap. The electron-sea model is a qualitative interpretation of band theory (molecular-orbital model for metals).

Molecular Orbital Theory Recall that atomic orbitals mix to give rise to molecular orbitals.

Molecular Orbital Theory In such elements, the energy gap between molecular orbitals essentially disappears, and continuous bands of energy states result.

Formation of Bands When atoms come together to form a compound, their atom orbital energies mix to form molecular orbital energies. As more atoms begin to mix and more molecular orbitals are formed, it is expected that many of these energy levels will start to be very close to, or even completely degenerate, in energy. These energy levels are then said to form bands of energy remember each orbital only holds two electrons.

Eventually you get some electrons on the conduction band in metals

The electronic band structure of nickel. The left side of the figure shows the electron configuration of a single Ni atom, while the right-hand side of the figure shows how these orbital energy levels broaden into energy bands in bulk nickel. The horizontal dashed gray line denotes the position of the Fermi Level, which separates the occupied molecular orbitals (shaded in blue) from the unoccupied molecular orbitals.

Solid State Materials- Metals and Metalloids

As metal and semi metal atoms bond to form a solid their bonding orbital combine into the molecular orbital model. 4 moles of atoms will generate 4 moles of orbitals. However each orbital can hold only two electrons.

Types of Materials Rather than having molecular orbitals separated by an energy gap, these substances have energy bands. The gap between bands determines whether a substance is a metal, a semiconductor, or an insulator.

Energy bands in metals, semiconductors, and insulators. Metals are characterized by the highest-energy electrons occupying a partially filled band. Semiconductors and insulators have an energy gap that separates the completely filled band (shaded in blue) and the empty band (unshaded), known as the band gap and represented by the symbol Eg. The filled band is called the valence band (VB), and the empty band is called the conduction band (CB). Semiconductors have a smaller band gap than insulators.

Metals Valence electrons are in a partially-filled band. There is virtually no energy needed for an electron to go from the lower, occupied part of the band to the higher, unoccupied part. This is how a metal conducts electricity.

Semiconductors Semiconductors have a gap between the valence band and conduction band of ~50-300 kJ/mol.

Semiconductors Among elements, only silicon, germanium and graphite (carbon), all of which have 4 valence electrons, are semiconductors. Inorganic semiconductors (like GaAs) tend to have an average of 4 valence electrons (3 for Ga, 5 for As).

Valence electrons in conduction band are release into the crystal Figure 12-00CO2 Title: HERE, THERE, AND EVERYWHERE. Caption: The element silicon is the primary component of computer processor chips and commercial solar panels. Notes: Keywords:

An intrinsic semiconductor is a semiconductor in its pure state An intrinsic semiconductor is a semiconductor in its pure state. For every electron that jumps into the conduction band, the missing electron will generate a hole that can move freely in the valence band. The number of holes will equal the number of electrons that have jumped. The higher the temp more electrons into conduction band.

Insulators The energy band gap in insulating materials is generally greater than ~350 kJ/mol. They are not conductive.

The relationship between orbital overlap and band gap (a). Figure 12-05a Title: Caption: Notes: Keywords: In diamond the C—C distance is relatively short (1.55 Å). This distance leads to effective overlap of orbitals on neighboring atoms, which in turn leads to a large splitting between the valence and conduction bands (Eg = 5.5 eV).

Longer the bond- weaker the bond-larger the gap In silicon the Si—Si distance is much longer (2.35 Å), which diminishes the orbital overlap leading to small splitting between the valence and conduction bands (Eg = 1.11 eV). The weaker the bond the less the gap.

The following pictures show the electron populations of the bands of MO energy levels for four different materials: (a) Classify each material as an insulator, a semiconductor, or a metal. Arrange the four materials in order of increasing electrical conductivity. Explain. Tell whether the conductivity of each material increases or decreases when the temperature increases. (b) (c)

Doping By introducing very small amounts of impurities that have more valence electrons (n-Type) or fewer (p-Type) valence electrons, one can increase or decrease the conductivity of a semiconductor.

The relationship between bond polarity and band gap. In germanium the bonding is purely covalent.

In gallium arsenide the difference in electronegativity introduces polarity into the bonds. The gallium atoms are less electronegative than germanium, which is reflected in an upward shift of the energies of the gallium atomic orbitals. The arsenic atoms are more electronegative than germanium, which is reflected in a downward shift of the energies of the arsenic atomic orbitals. The introduction of bond polarity increases the band gap from 0.67 eV for Ge to 1.43 eV for GaAs.

Presented By, Mark Langella, APSI Chemistry 2014 , PWISTA .com

The addition of controlled small amounts of impurities (doping) to a semiconductor changes the electronic properties of the material. Left: A pure, intrinsic semiconductor has a filled valence band and an empty conduction band (ideally). Middle: The addition of a dopant atom that has more valence electrons than the host atom adds electrons to the conduction band (i.e., phosphorus doped into silicon). The resulting material is an n-type semiconductor. Right: The addition of a dopant atom that has fewer valence electrons than the host atom leads to fewer electrons in the valence band or more holes in the valence band (i.e., aluminum doped into silicon). The resulting material is a p-type semiconductor. Figure 12-07 Title: Caption: Notes: Keywords:

Which of the following is a p-type semiconductor? Sulfur-doped carbon Boron-doped germanium Phosphorus-doped silicon Ultra-pure silicon Carbon-doped copper

Which of the following is a p-type semiconductor? Sulfur-doped carbon Boron-doped germanium Phosphorus-doped silicon Ultra-pure silicon Carbon-doped copper

Identifying Types of Semiconductors Which of the following elements, if doped into silicon, would yield an n-type semiconductor? Ga; As; C. Solution Solve: Si is in column 4A, and so has four valence electrons. Ga is in column 3A, and so has three valence electrons. As is in column 5A, and so has five valence electrons; C is in column 4A, and so has four valence electrons. Therefore, As, if doped into silicon, would yield an n-type semiconductor. Suggest an element that could be used to dope silicon to yield a p-type material. Answer: Because Si is in group 4A, we need to pick an element in group 3A. Boron and aluminum are both good choices—both are in group 3A. In the semiconductor industry boron and aluminum are commonly used dopants for silicon. Practice Exercise

Diode- Used to switch and convert between electromagnetic radiation and electric current Semiconductor created that has p-type on one half and n-type on the other half Known as “p-n rectifying junction” The energy level adjust so the Fermi Levels are equal This movement of levels creates valence level distortion and Conduction Band Distortion Valence Band n-type valence level is lower than p-type Conduction Band of of N-type is greater than p-type P- type has more “positive holes “ in Valence level Equilibrium Phenomena Holes from p-type valence band flow to n-type Fermi Level and electrons flow from n-type Valence Band into p-type Valence Band

Light Emitting Diodes p-type semiconductor “The addition of a dopant atom that has fewer valence electrons than the host atom leads to fewer electrons in the valence band or more holes in the valence band joined to a n-type semiconductor"the addition of a dopant atom that has more valence electrons than the host atom adds electrons to the conduction band “ When Voltage applied positive to p-type and negative “ electrons” to n-type When electron from conduction Band in n-type moves across junction it can drop into a “positive hole” in the valence band of p-type This drop emits energy based on the Gap difference between Conduction band of n-type and Valence Band of p-type. “ Band Gap”

Light emitting diodes. The heart of a light emitting diode is a p-n junction where an applied voltage drives electrons and holes to meet. Bottom: The color of light emitted depends upon the band gap of the semiconductor used to form the p-n junction. For display technology red, green, and blue are the most important colors because all other colors can be made by mixing these colors.

Color λ Voltage Drop Composition Red 610 < λ < 760 1.63 < ΔV < 2.03 Aluminium gallium arsenide (AlGaAs) Gallium arsenide phosphide (GaAsP) Aluminium gallium indium phosphide (AlGaInP) Gallium(III) phosphide (GaP) Orange 590 < λ < 610 2.03 < ΔV < 2.10 Gallium arsenide phosphide (GaAsP) Aluminium gallium indium phosphide (AlGaInP) Gallium(III) phosphide (GaP) Yellow 570 < λ < 590 2.10 < ΔV < 2.18 Green 500 < λ < 570 1.9[63] < ΔV < 4.0 Traditional green: Gallium(III) phosphide (GaP) Aluminium gallium indium phosphide (AlGaInP) Aluminium gallium phosphide (AlGaP) Pure green: Indium gallium nitride (InGaN) / Gallium(III) nitride (GaN) Blue 450 < λ < 500 2.48 < ΔV < 3.7 Zinc selenide (ZnSe) Indium gallium nitride (InGaN) Silicon carbide (SiC) as substrate Silicon (Si) as substrate—under development Violet 400 < λ < 450 2.76 < ΔV < 4.0 Indium gallium nitride (InGaN)

Relationship of Gap to Frequency The size of the band gap depends upon the vertical and horizontal positions of the elements in the periodic table. The band gap will increase when either of the following conditions is met: (1) The elements are located higher up in the periodic table, where enhanced orbital overlap leads to a larger splitting between bonding and antibonding orbitals: or (2) The horizontal separation between the elements increases, which leads to an increase in the electronegativity difference and the bond polarity.

Band Gaps and Wavelength

Qualitative Comparison of Semiconductor Band Gaps Will GaP have a larger or smaller band gap than ZnS? Will it have a larger or smaller band gap than GaN? Solution Plan: We must look at the periodic table and compare the relative positions of the elements in each case. Solve: Gallium is in the fourth period and group 3A. Its electron configuration is [Ar]3d104s24p1. Phosphorus is in the third period and group 5A. Its electron configuration is [Ne]3s23p3. Zinc and sulfur are in the same periods as gallium and phosphorus, respectively. However, zinc, in group 2B, is one element to the left of gallium and sulfur in group 5A, is one element to the right of phosphorus. Thus we would expect the electronegativity difference to be larger for ZnS, which should result in ZnS having a larger band gap than GaP. For both GaP and GaN the more electropositive element is gallium. So we need only compare the positions of the more electronegative elements, P and N. Nitrogen is located above phosphorus in group 5A. Therefore, based on increased orbital overlap, we would expect GaN to have a larger band gap than GaP. Additionally, nitrogen is more electronegative than phosphorus, which also should result in a larger band gap for GaP.

Qualitative Comparison of Semiconductor Band Gaps Will ZnSe have a larger or smaller band gap than ZnS? Answer: Because zinc is common to both compounds and selenium is below sulfur in the periodic table, the band gap of ZnSe will be smaller than ZnS. Practice Exercise

Presented By, Mark Langella, APSI Chemistry 2014 , PWISTA .com Figure 12.12 Figure 12-12 Title: Absorption of light at a p-n junction. Caption: This figure illustrates the process by which light is absorbed at the junction between a p-type and an n-type semiconductor. First, a photon is absorbed in the junction exciting an electron from the valence band into the conduction band creating an electronhole pair (this process is marked with a 1). Next the electron (e–) is attracted to the n-type semiconductor and the hole (h+) toward the p-type semiconductor (this process is marked with a 2). In this way the energy of the photon can be converted into electrical energy. Notes: Keywords: Presented By, Mark Langella, APSI Chemistry 2014 , PWISTA .com

Electricity from sunlight. Figure 12-13 Title: Caption: Solar panels, made of silicon, are used both as an energy source and as architectural elements in this apartment building in southern California. Notes: Keywords:

Structures of Metallic Solids Crystal structures of metals are simple enough we can generate the structure by placing a single atom on each lattice point Primitive cubic structure are rare – radioactive polonium Body-centered cubic metals – iron, chromium, sodium, and tungsten Face-centered cubic metals – aluminum, lead, copper, silver, and gold

CLOSE PACKING Shortage of valence electrons and the fact that they are collectively shared make it favorable for the atoms in a metal to pack together closely Atoms are spherical – understand structure of metals by looking at how spheres pack together Efficient way – pack one layer of equal sized spheres to surround each sphere by six neighbors For 3-D structure keeps stacking layer Second layer will sit in the depressions made by the first spheres

CLOSE PACKING Third layer – we have 2 choices First – third layer in the depressions that lie directly over the spheres in the first layer Fourth layer would lie directly over the spheres in the second layer (FIG 12-13) HEXAGONAL CLOSE PACKING (HCP)

CLOSE PACKING Second – third layer does not sit directly above the spheres in either of the first two layers, subsequent layers repeat this sequence giving an ABCABC pattern CUBIC CLOSE PACKING (CCP)

CLOSE PACKING In both each sphere has 12 equidistant nearest neighbors - 6 in the same layer 3 from the layer above and 3 from the layer below Each sphere has a coordinate number of 12 – the number of atoms immediately surrounding a given atom in a crystal structure

Ceramics These are inorganic solids, usually hard and brittle. They are highly resistant to heat, corrosion and wear. Ceramics do not deform under stress. They are much less dense than metals, and so are used in place of metals in many high-temperature applications. © 2009, Prentice-Hall, Inc.

Alloys Alloys contain more than one element and have the characteristic properties of metals. Solid Solution alloys are homogeneous mixtures. Pure metals and alloys have different physical properties. An alloy of gold and copper is used in jewelry (the alloy is harder than the relatively soft pure 24 karat gold). 14 karat gold is an alloy containing 58% gold.

Metal Alloys-Solid Solutions Substance has mixture of element and metallic properties. 1. Substitutional Alloy: some metal atoms replaced by others of similar size. Electronegativities usually are similar. The atoms must have similar atomic radii. The elements must have similar bonding characteristics. brass = Cu/Zn

Metal Alloys (continued) 2. Interstitial Alloy: Interstices (holes) in closest packed metal structure are occupied by small atoms. Solute atoms occupy interstices “small holes” between solvent atoms. One element (usually a nonmetal) must have a significantly smaller radius than the other (in order to fit into the interstitial site). steel = iron + carbon 3. Both types: Alloy steels contain a mix of substitutional (Cr, Mo) and interstitial (Carbon) alloys.

Alloys vs. Pure Metal The alloy is much harder, stronger, and less ductile than the pure metal (increased bonding between nonmetal and metal). An example is steel (contains up to 3% carbon). mild steels (<0.2% carbon) useful for chains, nails, etc. medium steels (0.2-0.6% carbon) useful for girders, rails, etc. high-carbon steels (0.6-1.5% carbon) used in cutlery, tools, springs. Other elements may also be added to make alloy steels. Addition of V and Cr increases the strength of the steel and improves its resistance to stress and corrosion. The most important iron alloy is stainless steel. It contains C, Cr (from ferrochrome, FeCr2), and Ni. Heterogeneous alloys: The components are not dispersed uniformly (e.g., pearlite steel has two phases: almost pure Fe and cementite, Fe3C).

Substitutional Alloy Interstitial Alloy

Ordered intermetallic structures. Keywords: (a) The face-centered cubic unit cell of a cubic close-packed metal. (b) The ordered structure of Ni3Al with nickel atoms shown in gray and aluminum atoms in blue. (c) The ordered structure of TiAl with titanium atoms shown in red and aluminum atoms in blue.

Figure 23.17abcd Figure 23-17 Title: Ordered intermetallic structures. Caption: (a) The face-centered cubic unit cell of a cubic close-packed metal. (b) The ordered structure of Ni3Al with nickel atoms shown in gray and aluminum atoms in blue. (c) The ordered structure of TiAl with titanium atoms shown in red and aluminum atoms in blue. (d) The ordered structure of ZnCu with copper atoms shown in copper and zinc atoms in gray. Notes: Keywords:

Which two substances are most likely to form an interstitial alloy? Nickel and titanium Silver and tin Tin and lead Copper and zinc Tungsten and carbon Solute atoms need to have much smaller bonding atomic radius than the solvent atoms typically a nonmetal that makes a covalent bond to the neighboring metal atoms – presence of the extra bonds provided by the interstitial component causes the metal lattice to become stronger and less ductile