Liquids, Solids & Modern

Liquids, Solids & Modern Materials: An Overview Phase Diagram: A graphical representation of the phases of matter • Critical point: Highest temperature & pressure required to have supercritical fluid. • Gas cannot be condensed to a liquid at this point no matter how high the pressure. • Triple point: Three phases of matter exist in equilibrium. • Gas, solid & liquid Reference: langlopress.net Supercritical Fluid: A 5th Phase of Matter • When a liquid is heated in a sealed container, a greater amount of vapor collects within the container causing the pressure inside to rise. • This results in the density of the vapor to increase and at the same time causing the density of the liquid to decrease. • At some temperature, the meniscus between the liquid and vapor disappears and the states commingle to form a supercritical fluid. • Supercritical fluids have properties of both gas and liquid states. Reference: allfordrugs.com Practice Problem: Reading a Phase Diagram Determine the phase(s) at each of the following points: a. 73.0 atm & (-56.4 oC) b. 5.11 atm & (-56.4 oC) c. 1 atm & (-78.5 oC) d. 73.0 atm & 31.1 oC e. Does solid CO2 float in liquid CO2? CO2 phase diagram Reference: chemistry.stackexchange.com Intermolecular Attractions (IM forces) • Intermolecular forces are attraction forces (i.e., electrostatic attractions) between molecules and atoms. • Kinetic molecular theory states the attractive forces between particles in a: • Gas are negligible; thus, gases have neither a definite shape or volume. • Liquids & solids have definite volumes, which means that there are attractive forces holding the particles together. • IM determine and influence a substance’s: • Molecular structure & phase it exists under STP conditions. • Phase reflects the strength (e.g., magnitude) of IM forces between molecules and/or atoms relative to thermal energy. • Physical properties (e.g., boiling point, melting point)of the substance. • The higher the boiling point of a liquid the stronger the intermolecular forces. • Boiling point is going from a liquid phase (moderate amount of intermolecular force interactions) to the vapor phase (no intermolecular interaction between molecules and atoms). Intermolecular Forces vs. Intramolecular Forces • Both are electrostatic interactions. • Intermolecular forces (IM) are electrostatic attractions between molecules and atoms. • Intramolecular forces are electrostatic attractions between two atoms to form a bond. Reference: openstax • Comparing one on one: intramolecular forces are stronger than intermolecular attractions. • Intermolecular forces have a cumulative effect on strength. • Have “many” intermolecular interactions in a solid or liquid not just one interaction. • Networking • Size and distance between molecules or atoms play a role in intermolecular attraction strength, similar to intramolecular forces. • Greater the distance the weaker the attraction Intramolecular vs. Intermolecular Forces: Types Intra molecular • Attractive forces that hold atoms together to form a bond. Keeps molecule intact. • Types: • Ionic • Metallic • Covalent • True • Polar Intermolecular • Attractive forces between molecules and atoms. • Types: • Ion-ion • Ion-dipole • Dipole-dipole • Hydrogen bonding • Dipole-induced dipole/ Ioninduced dipole • Induced-induced • Dispersion /Van der Waals Knowing the type of intramolecular forces (i.e., compound (ionic, molecular) and polarity of substance predicts the type of intermolecular forces present. Ion-ion: Strongest IM • Associated with ionic compounds (e.g. cation to anion interaction). • Crystalline structure • Ordered arrangement • Are strongest of intermolecular forces. • 400 to 4000 kJ/mole • Energy range due to charge and ion size. NaCl • Explains why ionic solids tend to have HIGH melting points temperatures. Reference: www.reddit.com Ion-dipole Interactions • Intermolecular attractions observed with a polar molecule (dipole) and ion. • Example: dissolution of ionic compound (e.g., salt) into water. • Strength of attraction is affected by the size of ion and polarity of strength of the dipole. • This interactions is 1/10th or 10% the strength of a covalent (intramolecular) attraction. • Energy: 40 to 600 kJ/mole Reference: www.ck12.org Dipole-Dipole Interactions • Occurs when molecules have a permanent dipole (e.g., polar molecules). • Polar molecules have a permanent dipole moment due to the following factors: • Bond polarity and molecular geometry (shape). • Strength of interaction varies: • Depends on polarizing strength. • Strength of these interactions are 1-5% the strength of a covalent bond. • Energy: 5 - 40 kJ/mole • Hydrogen bonding is a dipole-dipole interactions observed in molecules having O-H or N-H or F-H. Reference: www.slideshare.net Hydrogen Bonding: A Closer Look at H2O • H2O is a liquid at room temperature, whereas most molecular substances with similar molar masses are gases at room temperature (e.g., NH3, CH4). • This is due to the hydrogen-bonding between and among the H2O molecules . • Water is an excellent solvent for dissolving many ionic and polar molecular substances, because of its large dipole moment. • Even many small nonpolar molecules have some solubility in water (e.g., O2, CO2). • Physical properties of water: • Has a very high specific heat (4.184 J/goC)for a molecular mass • Water expands when it freezes so its solid (ice) less dense than liquid (water). Reference: witcombe.sbc.edu Induced Dipole Interactions: Ion-induced or Dipole-induced • The interaction between a “polar molecule or a charged particle (ion) with a nonpolar substance. • Two ways to induce an instantaneous ‘dipole moment’ in a nonpolar substance: • Ion or Dipole to induced dipole. • Strength of interaction varies from 2 to 10 kJ/mole. • Depends on: • Polarizability of the nonpolar substance. • Larger molecule (i.e., molar mass) more electron density then the greater the polarizability. • Charge and size of ion and the polarity of the polar molecule. • More surface area (i.e., shape of molecule or size of ion) the contact to induce. • This IM force explains how can nonpolar molecules such as O2 and I2 dissolve in water, a polar substance. Reference: faceweb.bhc.edu Dispersion forces: Induced-Induced Dipole • The interaction between nonpolar molecules that causes an instantaneous dipole (b) to form. • Factors that influence the strength of dispersion forces: • Polarizability of the nonpolar molecule, • Size of molecule (molecular mass). • Larger molecules more “interaction” sites. • The weakest of the intermolecular forces • 0.05 – 40 kJ/mole. • Also known as Van der Waal forces. Symmetric distribution of electron density Asymmetric distribution of electron density Reference: chemistry.elmhurst.edu Dispersion Forces Factors: Molecular Shape & Molecular Mass • More “contact” spots the greater the effect dispersion forces have on a molecule's physical properties such as boiling point. • Example: • n-pentane (36 oC) vs. Neopentane (2,2 – dimethyl propane )(10 oC) Reference: Openstax Problem: Predict whether any of the following pairs of compounds can form ion– dipole attractions. a. H2O and CO2 b. NaCl and H2O c. NaCl and CO2 Problem: Predict the dominant type of intermolecular force between the following: a. H2O & CO2 b. NaCl & H2O c. CH2O d. CCl4 e. HF f. NH4+ g. H2 h. PCl3 Intermolecular Forces & the Physical Properties of Liquids • Surface tension • Capillary Action • Viscosity • Solubility • Boiling points Surface Tension Reference: ouitschem.weebly.com • A liquid’s resistance and associated energy require to increase its surface area. • Liquids spontaneously seek to minimize their surface area. • The energy required is directly related to the type of intermolecular forces within the liquid. • Stronger the intermolecular forces, the larger amount of energy is required to increase the liquid’s surface area. • Examples of surface tension energies at room temperature: • H2O = 72.8 mJ/m2 • C6H6 = 28 mJ/m2 • Temperature affects a liquid’s surface tension. • Increasing the temperature, reduces the surface tension energy. • With an increase in molecular kinetic energy, have more energy to break the intermolecular attractions. Capillary Action: • The spontaneous rising of a liquid in a narrow tube or the movement up a piece of paper against the pull of gravity. • There are two ‘forces’ at play in capillary action of a liquid: • Cohesive forces: the forces associate within the liquid. • Adhesive forces are the attractive forces that the liquid’s outermost molecules have with the paper or tube surface. • Shape of liquid’s meniscus is an indication whether similar polarities exist between the liquid and the container. • Concave: adhesion forces between the liquid and the composition of the material of the container is stronger than its own cohesion forces. • Convex: the adhesion forces to the container wall is less than the liquid’s cohesion forces Reference: vison.eecs.ucf.edu Viscosity • Viscosity is a measurement of a liquid’s resistance to flow, and its associated energy required to move an object through a liquid. Reference: opentextbc.ca • Factors that affect the viscosity of a liquid: • Temperature: • Increase in temperature (i.e., more kinetic energy) the liquid’s viscosity decreases. • Size of liquid’s molecule: • The more complex, greater entanglement; larger molecule more viscous. • Intermolecular forces of the liquid. Intermolecular Forces & the Solubility Connection • Solubility depends, in part, on the attractive forces of the solute and solvent molecules. • Similar polarities or “like dissolves like.” • Miscible liquids will always dissolve in each other. • Polar substances dissolve in polar solvents. • They have hydrophilic groups (water loving) • Nonpolar substances dissolve in nonpolar solvents. • Their compounds have hydrophobic groups (fear of water). • Immiscible liquids do not mix; they have NOTHING in COMMON with each other. – Do not have anything in common (i.e., compatible IM attractions) that allows them to interact with or attract to each other. Intermolecular Forces & Boiling Point of a Liquid • Intermolecular forces affect a liquid’s boiling point. • The stronger the intermolecular forces between liquid particles the higher its boiling point. • Boiling point of a liquid is when the temperature of the liquid reaches a temperature where the liquid’s vapor pressure is equal to the external pressure. • It is where vapor (e.g., bubbles) can form anywhere within the liquid, not just on the surface. • The normal boiling point of a liquid is the temperature at which the liquid’s vapor pressure equals the atmospheric pressure. • The lower the external pressure, the lower the boiling point of the liquid. • Altitude increase → Atmospheric pressure decreases → Boiling point decreases Reference: infoplease.com Kinetic Molecular Theory of Liquids: Overview • The intermolecular forces present in a pure liquid (i.e., solvent only) are appreciable . • Polar substances : Dipole- Dipole • Nonpolar substances: Dispersion • Liquid particles have close contact with one another which limits their movement but does not hinder their ability to flow. • Liquids have higher densities than gases and are incompressible because the particles are in contact. • A liquid’s volume is the capacity of the container. • They have an indefinite shape because the limited translational freedom of the particles allows them to move around enough to get to the container walls. • However, they have a definite volume because the limit on their freedom keeps them from escaping the rest of the particles. Vaporization: Liquid to Gas • Vaporization is an endothermic process ; requires the input of energy to overcome the IM attractions between the liquid particles. • Increasing the temperature increases the kinetic energy of the liquid molecules resulting in the fraction of the molecules with “escape energy” to increase. • The weaker the attractive forces between molecules, the less energy to vaporize and the faster the rate of evaporation. • The greater the liquid’s surface area, the faster the rate of evaporation. • Liquids that evaporate easily are said to be volatile. • e.g., gasoline, fingernail polish remover • Liquids that do not evaporate easily are called nonvolatile. • e.g., motor oil, salt water Reference: hk-phy.org Condensation: Gas to Liquid • Condensation is an exothermic process. • It is the opposite of vaporization. • Molecules in the vapor state can lose energy through molecular collisions with other molecules. • The result: • Molecules get captured back into the liquid when they collide with the liquid’s surface or may stick together to form droplets of liquid. • The weaker the intermolecular attractive forces, the more energy is needed for the vapor (e.g., gas) particles to condense. Reference: processtechacademy.com Phase Diagrams & Phase Changes (DH) • Phase changes are thermodynamic processes. • Exothermic or Endothermic • The strength of the attractive forces (intermolecular) between particles determine the amount of energy needed or released when changing states. Reference: kentchemistry.com DHvap and Intermolecular Connection • DHvap is the amount of energy required to vaporize one mole of liquid to the vapor phase. It is temperature dependent. • LIQUID + heat → VAPOR • + DHvaporization = - DHcondensation • The vapor pressure of a liquid at a given temperature depends on the liquid’s intermolecular forces. • Stronger and/or more networking of IM forces - the higher the vapor pressure. • An increase in mass results in more IM forces present, thus the greater the vapor pressure. Compd. ∆Hvap (kJ/mol) Intermolecular Force (IM) H2O 40.7 (100 oC) H-bonds SO2 26.8 (-47 oC) Dipole-dipole Xe 12.6 (-107 oC) induced dipole–induced dipole Heating Curves • As a substance is heated at a constant rate, the substance’s temperature will increase linearly until it reaches its boiling point. • At each of the transition points (e.g., melting point) the temperature of the substance remains constant. • All the added energy goes into melting the solid; that is why the temperature remains constant. • The temperature remains constant until the next transition phase at which point the temperature will again start to rise. Reference: scienceuniverse101.blogspot.com Problem: Determine the amount of energy needed to evaporate 100.0 grams of liquid ammonia to its boiling point. DHvap = 4.8 kJ/mol Equilibrium Vapor Pressure • The equilibrium vapor pressure is defined as the pressure exerted by a vapor over a liquid in a closed container when the rate of evaporation = the rate of condensation. • When molecules of liquid are in the vapor state, they exert a VAPOR PRESSURE on the surface of the liquid. • Evaporation and condensation are still occurring, but because they are opposite processes, there is no net gain or loss of either vapor or liquid. Reference: www.kmacgill.com Vapor Pressure & Intermolecular Connection Vapor pressure curves • The stronger the attractions between the atoms or molecules, the more energy it will take to go from one phase transition to another. • Boiling a liquid requires addition of enough energy to overcome all the INTERMOLECULAR attractions between the particles. • NOTE: INTRAMOLECULAR (e.g., covalent, ionic, metallic) forces DO NOT BREAK. • The higher boiling point of the liquid, the stronger the intermolecular attractive forces and the greater the vapor pressure of the liquid. Reference: www.800mainstreet.com Vapor Pressure Curves • Vapor pressure versus temperature curves graphically represent the relationship between changes in temperature and pressure. Vapor pressure curves • As the temperature increases the number of molecules able to escape the liquid will also increase. • The net result is that as the temperature increases, the vapor pressure increases. • Small changes in temperature make large changes in the vapor pressure. • The magnitude of the changes depend on strength of the intermolecular forces present in the liquid. Reference: www.800mainstreet.com Clausius-Clapeyron Equation • The Clausius-Clapeyron equation is the linear relationship used to find ∆H˚vap when the vapor pressure and temperature is known. • The logarithm of the vapor pressure P is proportional to ∆Hvap and to 1/T. • The linear relationship between vapor pressure its temperature is: ln P = (-∆H˚vap /R) (1/T) + Constant • This equation can be used with just two measurements of vapor pressure and temperature. ln (P2/P1) = {-(DHvap/R) [(1/T2) –(1/T1)]} Clausius-Clapeyron Relationship • The natural logarithm of the vapor pressure (ln Pvap) versus inverse absolute temperature is a linear function. • A graph of ln Pvap versus 1/T is a straight line. • The slope of the line: - (ΔHvap / R) • Where R is 8.314 J/mol ∙ K Reference: intro.chem.okstate.edu Question? b Which line (a) or (b) has: 1. The smaller DHvap a 2. The stronger intermolecular forces? Problem #1: Determine the DHvap for propyl amine from the following data. P1 = 40.0 mm Hg T1 = 257 K P2 = 100.0 mm Hg T2 = 274 K R = 8.314 x 10-3 kJ/K mol ln (P2/P1) = {-(DHvap/R) [(1/T2) –(1/T1)]} Problem #2: Determine the P2 in mm Hg from the following data. P1 = 400.0 mm Hg T1 = 2.59 x 103 K R = 8.314 x 10-3 kJ/K mol P2 = ? mm Hg T2 = 2.62 x 103 K DHvap = 252 kJ/mol ln (P2/P1) = {-(DHvap/R) [(1/T2) –(1/T1)]} Please Note: the some of the following material on Material Science - is not found in your textbook. Material Science: The Study of Solid Material Properties • Material Sciences, specifically investigate how a solid’s composition and structure (i.e., atomic and molecular) affect its chemical and physical properties. • Understanding of properties leads to application possibilities. • It is an amalgam of the science disciplines of Chemistry, Metallurgy, Physics, Engineering and Design. • Divisions of material sciences: metals (i.e., metallurgy), polymers, semiconductors, and ceramics. Grouping of Material Science Solids • Main groupings are defined according to atomic bonds and structure: • Metals or metallic solids • Ceramics: ionic solids or covalent network solids • Polymers: covalent molecular solids • Composites: combination of two or more of the solid types (i.e., metals, ceramics or polymers. • NOTE: Classification of solid materials can also be according to: • Specific physical properties (e.g., mechanical, electrical, optical, etc.). • Application or use of the solid material (e.g., structural component, device, machine, etc.). Overview of Solid Properties • Particle composition: The atoms, ions and/or molecules are tightly packed together which result in solids being incompressible, having high densities & defined volumes. • Particle motion: Solids have little, if any translational or rotational freedom of motion. • Explains why solids retain their shape and volume when placed in a new container; they is no flow of particles. • Particle order: The arrangement of particles within the solid exist. • Types/categories: • Crystalline • Polycrystalline • Amorphous: Kinetic Theory of Solids * The Kinetic Theory of Matter states that all matter forms have motion (movement). It is the level of motion (degrees of freedom) that defines the phase. • Kinetic Theory of Solid: • When the attractive forces between particles (i.e., molecules or atoms) are strong enough so the kinetic energy cannot overcome it at all, the material will be a solid. • The forces between the solid particles depend on the type of particles. • Attractive forces can be intermolecular (ion-ion, dipole, dispersion etc.), intramolecular (covalent, metallic or ionic) or a combination of inter and intra molecular forcers. • The packing of solid particles is such that only vibrational motion (a movement of freedom) exists; very little, if any translational or rotational motion is observed. Particle Motion – Degrees of Freedom • There are three degrees of freedom (i.e., movement) a particle can possess: translational, rotational & vibrational. • All phases of matter possess vibrational motion. • Temperature influences particle motion. • Vibrational motion has two basic groupings: • Bending – change of bond angle. • Four changes: scissoring, rocking, wagging and twisting • Stretching – continuous change of movement within and along the axis of a bond. • Spectral analysis takes advantage of particle motion. Reference: slideplayer Crystalline: • Orderly geometric pattern of same size particles and orientation through out solid. • Solid particles are bonded in a constant & coordinated manner. • Bond angles and distances between gas particles describe structure type. • Examples: diamond, salt Polycrystalline: • Semi ordered or intermediate crystal pattern. • Composed of varying size particles and crystalline regions having different orientation patterns with in the solid. • These regions are referred to as grains. • Examples: Ice, pure metals, ceramics Amorphous: • Non regular geometric pattern over a long range. • Solid particles tend to form short-range area of periodic structure. • Example: plastic and glass Reference: Top pr com. & studylib.net Problems: 1. Which of the following represents lattice energy? A. 2Al(s) + 3 Cl2(g)  2 AlCl3(s) B. Al3+(g) + 3 Cl–(g)  AlCl3(s) C. Cl(g) + e –  Cl–(g) 2. Which of the following ionic solids would have the greater lattice energy. A. KF B. LiBr C. NaCl Silicates • Silicates compose ~90% of Earth’s crust. • Glass & Quartz (sand) are types of silicate compounds. • Glass is the amorphous form. • Quartz is the crystalline version. • NOTE: SiO2 in pure form is clear but impurities (metal atoms) adds color. • Silicates have a three-dimensional array of Silicon atoms covalently bonded to four Oxygen atoms. Reference: sas.upenn.edu Glass: A Randomly Amorphous Solid Glass is formed by quickly cooling liquid silica (SiO2 (sand). • Historically, found in many ancient cultures (e.g., Greeks, Egyptians, Romans), but the Romans discovered glass making and blowing as a technique. • Their discovery: Adding of Na2CO3 to melted silica allowed glass to form at lower more manageable temperatures. • Types of glass: • Vitreous or fused silica: pure silica liquefied then cooled quickly. • Modern, Window or soda-lime glass: 70% silica with remainder CaO or Na2O. • Borosilicate (Pyrex@): Soda-lime glass mixed with boric acid (B2O3) added not CaO. • This combination allows for thermal expansion, lessens thermal shock seen in soda-lime glass. • Leaded glass or crystal: When PbO is mixed with liquefied silica, resulting cooled substance structure results in a high index of refraction giving glass a bright brilliant look. • Storing liquids/drinks for prolong period of time in leaded glass will result in some leaching of the lead. Ceramics • Composed of inorganic nonmetallic powdered solids that when mixed with water can be molded and then heated to make items such as bricks, pottery, tiles and dishware. • Ceramic materials are nonconductive, strong, hard but brittle. • Types of ceramics: • Silicate ceramics are made up of aluminosilicates, which when weathered form clays. • Kaloinite, (Al2Si2O5(OH)4), when fired (heated above 1500 oC) changes both chemically and structurally to produce porcelain (white ceramic). • Oxide Ceramics, (MgO & Al2O3), when compared to their metal content have much higher melting points (3x) but are brittle and subject to thermal shock (e.g., crack with sudden changes in temperature). • Nonoxide Ceramics, ( BN, SiC, Si3N4) are network covalent solids that whose atoms are linked together in tetrahedral matrix. Cement • A binding material: • Binding material is substance that hardens independently while “binding” other materials together. • A fine, powdery material. • Factory produced CaO (60-67%), SiO2 (17-25%), Al2O3 (3-8% & Fe2O3 (0.5-6%) • Historically: discovered by the Romans when they mixed lime (Ca(OH)2) with volcanic ash and clay resulting in a slurry that formed when dry a hard, durable solid. • Portland cement (patented by a British stonemason, Joseph Aspdin (1824)). • Called Portland cement because when hardened it resembles a building stone found on the Isle of Portland in southwest England. • Composition: limestone (CaCO3), sand (SiO2) with small amounts of alumina (Al2O3), iron (III) oxide and gypsum (CaSO4 2H2O) and water. • As the cement hardens Si-O-Si linkages (e.g., bridges) form a fibrous like structure that hard, strong and durable. • If pebbles are added to Portland cement mixture – it is referred to as concrete. Crystalline Solids: Categories Crystalline solids classifications can be based upon: 1. Its particle make-up or 2. The types of attractive forces holding the particles together. Reference: askIITians Overview of Crystal Structure & Unit Cells • Crystalline solids are orderly geometric structures whose lattice points are occupied by atoms, ions, or molecules. • Crystal Lattice: three-dimensional arrangement of particles which result in achieving maximum attractive forces between particles in solid. • The number of other particles each particle is in contact with is known as its coordination number. • For ions, it is the number of oppositely charged ions an ion is in contact. • Higher coordination number: • more interaction; stronger attractive forces hold the crystal together. • Packing is more efficient per volume of unit cell. Unit Cells Categories Reference: epionelynx.wordpress.com Atomic Solids • Solids whose composite particles are at the lattice points of the crystal are atoms. • The intermolecular forces present in solid are dictated by atom type. • Nonbonding atomic solids are held together by dispersion forces. • These solids tend to have low melting points • Metallic atomic solids are held together by metallic bonds. • Network covalent atomic solids are held together by covalent bonds – polar covalent or covalent. • Due to the directionality of the covalent bonds, these solids do not tend to form closest-packed arrangements in the crystal. • The strength of the covalent bonds within the solid is why these solids tend to have very high melting points. • Generally, >1000 °C Allotropes of Sulfur Reference: chemistry libretexts Allotropes of Carbon Solids • Carbon solids have either a two-dimensional or three-dimensional structure. • Example: diamond is three-dimensional (tetrahedral, sp3) and graphite is two-dimensional (trigonal planar, hybridization of sp2. • Each carbon atom is surrounded by four covalently bonded carbon atoms. • This effectively makes each crystal one giant molecule held together by covalent bonds. Reference: archive.fabacademy.org Molecular Solids • The lattice sites of the crystal are occupied by molecules, such as CO2, H2O, C12H22O11 • The molecules are held together by intermolecular attractive forces. • Dispersion forces, dipole–dipole attractions, and Hydrogen bonds. • Because the attractive forces are weak, they tend to have low melting points. • Generally, <300 °C. DRY ICE DRY ICE ICE Reference: Commons.wikimedia.org Reference: Chemguide Reference: tutorMyself chemistry Ionic Solids • Composite particles are ions (e.g., cations or anions) and whose lattice sites in the crystal are occupied by ions. • Arrangement is in a repetitive three-dimensional pattern (unit cell). • Unit cell packing dictated by size and charge of ions present. • Attractive forces present ion-ion interactions. • Properties: hard, brittle, high melting points, • The ion-ion interactions explain very high melting points of solids. • Only when in dissolved in water (ionized) or molten state will ionic salts conduct electricity. Reason: ions of compound can move. Reference: wps.prenhall.com Metallic Solids • Metallic element atoms held together by metallic bonds. • The strength of these columbic interactions (attractions) varies with sizes and charges of metallic cations. • Arrangement in crystal, mostly closed-packed unit cells. • Properties: • Malleable (i.e., can be formed in different shapes), ductile (i.e., can be drawn/pulled to form wire), conduct heat & electricity, wide range of melting points • Metal atoms within solid release valance electrons • Metal atom cation is “fixed in a sea of mobile electrons.” Reference: Socratic Metallic Bonding & Band Theory • Qualitative description of metallic bonding is the “sea of electrons.” • Quantitative description of metallic bonding is Band Theory Model • Uses concept of bonding as described by Molecular Orbital Theory (MO) for molecular substances. • Electrons in atomic orbitals enter molecular orbitals – bonding & antibonding. • The area of bonding molecular orbitals is called the valence band. • The area of antibonding molecular orbitals (i.e., high energy orbitals) is called the conduction band. • The difference in energy between the valence band and conduction band is called the band gap. • Band gap, in a solid, is the is an energy range where electron states do not can exist. • The larger the band gap, the fewer electrons there are with enough energy to make the jump. • As the temperature rises, some of the electrons may acquire enough energy to jump to the conduction band. • Band gap (also known as Femi level) is the distance between the electrons in the valance band & conduction band. • Gap represents the MINIMUM energy need to excite an electron from on band level to the other. Reference: Quora & Energy Education Conductivity & Band Gap Relationship The more electrons occupying/residing within the conduction band the better it is as a conductor. • Conductor: Overlapping or has Little or no band gap. • Example: metals • The conductivity of a metal decreases with temperature. • Semiconductor: Very small band gap • At room temperatures, there will be a significant number of the electrons in the conduction band at normal. • Conductivity is temperature dependent; with increasing temperatures more electrons within the conduction band. • Example: graphite • The conductivity of a semiconductor increases with temperature. • Insulator: Very large band gap • Basically, no electrons will be in the conduction band at normal temperatures Ways to increase electron movement to conduction band: Doping • Doping: the addition of impurities to the semiconductor’s crystal to increase its conductivity. • n-Type semiconductors do not have enough electrons themselves to add to the conduction band, so they are doped by adding electron-rich impurities. • Adding an extra level of electron density • p-Type semiconductors are doped with an electron-deficient impurity, resulting in electron “holes” in the valence band. • Electrons can jump between these holes in the valence band, allowing conduction of electricity. • p-n junction: melding of p-type with a n-type semiconductor. • Diode: when electricity flows in one direction across a junction. • Amplifier: increases electrical current by allowing the accumulation of electrical energy Reference:chem.libretexts.org Polymers • Polymers are very large molecules made by repeated linking together of small molecules referred to as monomers. • Polymerization is the process of linking the monomer units together. • Two methods: • addition polymerization • condensation polymerization • Monomer units can be linked head to tail (most common) , or head-to-head, or tail to tail during polymerization. • Structural pattern of monomer units and attractions between strands affect polymer properties. Reference: Chemistry LibreTexts Reference: The Essential chemical Industry Polymers Classification • Polymers can be grouped according to their: • Structure: linear, cyclic, branched or networked. Reference: pocketdentisry.com • Source: • Natural polymers are polymers found in both the living and nonliving environments. • Synthetic polymers are polymers made in a lab from one, two or three small molecules linked in a repeating pattern. • Examples: Plastics, elastomers (rubber), fabrics, adhesives. • Composites polymers are materials made of polymers mixed with various additives. • Additives such as graphite, glass, metallic flakes. Polymers Classification • Polymers can also be grouped according to their: • Molecular interactions: thermoplastic, thermosetting, elastomers, fibers Semi-crystalline structure Thermoplastics: Can be melted and reshaped due to weak interactions between chains Amorphous structure Synthetic Polymer Thermosets: Cannot be melted. Polymerization resulted in rigid, cross-linked structured formed which heating cannot break in order to reshape. Green plastics • Green plastics are polymers derived from renewable sources biomasses, such as plants and animals, not oil or petroleumbased resources. • There are: • Biodegradable • Recyclable process are friendly to environment. Reference: Wikipedia