Chem 16 Advance Placement Exam PDF

University of the Philippines Chemical Engineering Society, Inc. (UP KEM) Chem 16 Advance Placement Exam Module Jardiolin © UP KEM 2324  University of the Philippines Chemical Engineering Society, Inc. (UP KEM) Chem 16 APE Module...

University of the Philippines Chemical Engineering Society, Inc. (UP KEM) Chem 16 Advance Placement Exam Module Jardiolin © UP KEM 2324  University of the Philippines Chemical Engineering Society, Inc. (UP KEM) Chem 16 APE Module Physical change - when a substance alters its physical form, Outline not its composition (e.g. phase change, cutting an object into pieces) 1. Matter: Properties, Composition, and the Changes it Undergoes Chemical change 2. Quantum Theory and Atomic Structure - when a substance is converted into a 3. Chemical Periodicity different substance (e.g. combustion, 4. Chemical Bonding oxidation, hydrolysis) 5. Molecular Geometry and Bonding Theories - One or more substances are used up. 6. Stoichiometry and Chemical Formulae - One or more new substances are formed. 7. Stoichiometry of Chemical Reactions This is accompanied by change in color 8. Thermochemistry 9. Liquids, Solids, Phase Changes and/or formation of precipitates or 10. Gasses bubbles. 11. Solutions - Energy is absorbed or released. 12. Acids and Bases 13. Changes in the Nucleus States of Matter MATTER: PROPERTIES, COMPOSITION, AND Solid THE CHANGES IT UNDERGOES - rigid, has definite shape, volumes do not vary that much with changes in temperature Chemistry and pressure - describes matter (its properties, the Liquid changes it undergoes, the energy that - flows, assumes the volume of the container, accompanies those changes). volumes do not vary that much with Matter changes in temperature and pressure - anything that has mass and occupies space Gas Energy - fills any container completely, expands - the capacity to do work or to transfer heat infinitely, easily compressed Potential energy - energy due to the position of the object or New States of Matter energy from a chemical reaction Kinetic energy Plasma - energy due to the motion of the object - Forms when atoms are heated more, they Potential and kinetic energy can be interconverted. split into free ions and electrons Liquid-Crystal - Transition phase between completely liquid Properties of Matter and completely solid phase - Has liquid fluidity but regular arrangements Physical property as with solids - exhibited by a substance by itself, without Bose-Einstein Condensate changing into or interacting with another - At absolute zero K conditions substance (e.g. color, melting point, electrical conductivity, density) - Extensive Property Measurement of Matter dependent on the amount of the substance present (volume and Accuracy mass) - closeness of the obtained measurement to - Intensive Property the true value independent of the amount of the - low systematic error substance present (color and Precision melting point) - closeness of obtained measurements to Chemical property each other - exhibited by a substance as it changes into - reproducibility of the measurement or interacts with another substance (e.g - low random error flammability, corrosiveness, reactivity with acids) 1 | UP KEM 2324 © KMCJ  University of the Philippines Chemical Engineering Society, Inc. (UP KEM) Chem 16 APE Module Length - the distance from one point to another Volume - the space occupied by matter Mass - the measure of the quantity of matter a body contains SIgnificant Figures Density 1. All non-zero numbers are significant. - an intensive property that relates two 452.58 → 5 significant digits extensive properties: mass and volume 2. Leading zeros are never significant. Percent Composition 0.000012 → 2 significant digits - Percent of a component in relation to the 3. Trailing zeroes may be significant total mass of the compound - specify significance by how the - %comp of x = [mass X/ total mass] x100% number is written Specific Gravity = (density of substance/density of 120 apples → indeterminate number of water) [unitless] significant digits since it is unknown Heat whether it was counted or weighed - a form of energy that always flows 6.245 × 105 → 4 significant digits spontaneously from a hotter body to a 4. Imbedded zeroes are always significant. colder body 703.20503 → 8 significant digits Temperature - measure of the intensity of heat Significant Figures on Addition/Subtraction Kelvin (K) is the “absolute - the number of decimal places in the temperature scale” begins at sum/difference should be equal to the absolute 0 and only has positive lowest number of decimal places (D.P.) in the values addends/minuend and subtrahend Conversion of Temperature 52.167 – 24.4 = 27.767 = 27.8 T (in K) = T (in °C) + 273.15 (3 D.P.) (1 D.P.) (1 D.P.) T (in °C) = T (in K) - 273.15 T (in °F) = (9/5) T (in °C) + 32 78.65 + 8.98673 = 87.63673 = 87.64 T (in °C) = [T (in °F) – 32] (5/9) (2 D.P.) (5 D.P.) (2 D.P.) Significant Figures on Multiplication/Division Composition of Matter - the number of significant figures in the product/quotient should be equal to the lowest number of significant figures (S.F.) in the factors/dividend and divisor 1.0000 × 2.5 = 2.5 (5 SF) (2 SF) (2 SF) 8.4000 ÷ 0.500 = 16.8 Pure Substances (5 SF) (3 SF) (3 SF) - fixed composition - cannot be separated into simpler Measurements of Matter substances by physical methods - can only be changed in identity and properties by chemical methods Physical Property Name of Unit Symbol - properties do not vary Length meter m Elements Mass kilogram kg - cannot be decomposed into simpler Time second s substances by chemical changes Electric current ampere A Compounds Temperature kelvin K - can be decomposed into simpler Luminous intensity candela cd substances by chemical changes, always in a definite ratio. Amount of mole mol Mixtures substance - variable composition 2 | UP KEM 2324 © KMCJ  University of the Philippines Chemical Engineering Society, Inc. (UP KEM) Chem 16 APE Module - components retain their characteristic - Discovered the nucleus and proton properties and proposed the existence of a - may be separated into pure substances by nucleus physical methods - Millikan: Oil drop Experiment - mixtures of different compositions may have - Discovered the mass of an electron widely different properties in 1909 Homogeneous Mixtures - have same composition throughout - components are indistinguishable Atomic Symbol and Isotopes of the Elements Heterogeneous Mixtures - do not have the same composition throughout - components are distinguishable Law of Conservation of Mass - Antoine Lavoisier - the total amount of matter (mass) remains constant regardless of any chemical change; matter cannot be created nor destroyed. Law of Definite Composition Where, - Joseph Louis Proust Z - atomic number, # of protons (also # of - pure samples of a compound will be electrons if neutral atom); constant for all composed of the same elements with atoms of the same element constant mass fractions A - mass number; may vary per element Law of Multiple Proportions Isotopes: same Z, different A - John Dalton Isobars: atoms of different elements - when two elements, A and B, form more than having the same atomic mass but one compound, the ratio of the masses of B different atomic number that combine with a given mass of A in each Isotones: elements having the same compounds can be expressed as small number of neutrons whole numbers And Law of Conservation of Energy number of neutrons = A - Z - Energy cannot be created or destroyed in a number of electrons = Z - charge chemical reaction or in a physical change. atomic weight (given % abundance of It can only be converted from one form to isotopes) = (%abundance of Isotope A)(Mass another. of Isotope A) + (%abundance of Isotope Dalton's Atomic Theory B)(Mass of Isotope B) +... - an element is composed of extremely small Atom indivisible particles called atoms. - Smallest particle of an element that - atoms cannot be created, destroyed, or maintains its chemical identity through all transformed into atoms of another element. physical and chemical changes - all atoms of a given element have identical Molecule properties (e.g. mass), which differ from - Smallest particle an element or compound those of other elements. that can have a stable independent - compounds are formed when atoms of existence different elements combine with each other Ions in small whole number ratios - When an atom or molecule gains or loses - Relative numbers and kinds of atoms are multiple or single electrons, they gain extra constant in a given compound charge that makes their properties entirely Atomic Models different from the original atom or - Democritus: Hard Sphere Model molecules - J.J. Thompson: Plum-pudding model Fundamental Particles - Discovered electrons in 1897 - Electrons (e-) = 0.00054858 amu - Used cathode ray tubes - Protons (p+) = 1.0073 amu - Neils Bohr: Planetary Model - Neutrons (no)= 1.0087 amu - Schrodinger: Electron Cloud Model - 1 amu = 1.66054 x 10-24 g - Rutherford: Gold Foil Experiment - Prior to his gold foil experiment, he first discovered α, β, and γ radiation 3 | UP KEM 2324 © KMCJ  University of the Philippines Chemical Engineering Society, Inc. (UP KEM) Chem 16 APE Module 2. Electrons moves in a circular orbit about QUANTUM THEORY AND ATOMIC the nucleus STRUCTURE 3. Electron may move from one discrete energy level (orbit) to another Atomic Spectra - light from excited gaseous atoms passing through a slit and refracted by a prism produces a line spectrum unique to that element - contradicted classical theory (Rutherford’s atomic model): electrons in an atom were expected to emit radiation with changing frequency, producing a continuous Basic Properties of Light spectrum Bohr's Model - line spectra of H atom is explained with this Wave Nature of Light - each atom has only certain allowable - relationship between frequency, v (number energy levels (stationary states) associated of cycles per second) and wavelength, λ with a fixed circular orbit of the electron (distance traveled by wave per second): around the nucleus - the atom does not radiate energy while in c=v*λ one of its stationary states (even if the where electron is moving within its orbit). c = 3.00 1O^ m/s - the atom changes to another stationary v = frequency or the number of state (the electron moves to another orbit) crests or troughs that pass a given only by absorbing or emitting a photon point per second whose energy equals the difference in - frequency and wavelength are inversely energy between the two states: proportional; the higher the frequency, the Ephoton = Estate A - Estate B =hv shorter the wavelength - amplitude (height of the crest of each wave) Quantum Model of an Atom describes the intensity of the light 1. Atoms can exist only in certain energy - evidence: diffraction pattern produced states. when light is passed through slits 2. Atoms emit or absorb radiation/light as gamma < x-ray < UV < violet < blue < green < yellow < they change their energies orange < red < infrared < microwave < radio 3. The allowed energy levels of electrons in an - arranged in inc. wavelength (dec. frequency) atom can be described by sets if number - visible region of light: ROYGBV called quantum numbers The 3 Phenomena that the wave model can’t explain Werner Heisenberg’s Uncertainty Principle 1. Emission of light from hot objects or - It is impossible to simultaneously determine black-body radiation the position and momentum of an electron 2. Emission of electrons from metal surfaces (or any other small particle) on which light shines or the photoelectric Erwin Schrodinger’s Schrodinger Equation effect - This equation gives us the wave function, or 3. Emission of light from electronically excited obitals or states of electron gas atoms to form emission spectra - The square of the equation is the Black-Body Radiation probability density - When solids are heated, they emit radiation Particle Nature of Light - energy of an atom is quantized (in discrete Quantum Numbers “packets", not continuous) ∆Eatom=Eemitted/absorbed radiation =∆nhv Quantum Numbers Where - Solutions to the Schrodinger, Heisenberg n is an integer and Dirac equations h = 6.626 1O-34 Js (Planck's constant) principal (n) - evidence: blackbody radiation, - orbital energy (size) photoelectric effect - positive integers (1, 2, 3,...) Atomic Spectra and Bohr Model angular momentum (I) 1. An atom has a number of definite and - orbital shape discrete energy levels or orbitals - integers from 0 to n —1 4 | UP KEM 2324 © KMCJ  University of the Philippines Chemical Engineering Society, Inc. (UP KEM) Chem 16 APE Module magnetic (ml) Atomic Radii - orbital orientation - generally, as size increases, reactivity - integers from -1 to 1 increases; the valence electrons are farther spin (ms) from the nucleus and are easier to remove - direction of e spin - shielding effect: atomic radii decreases - +1/2 or -1/2 through a row because the pulling becomes - each orbital can hold up to two electrons, stronger as the number of protons increase and they must have opposite spins (m,) Ionization Energy Pauli's Exclusion Principle - first ionization energy: minimum energy - no two electrons of the same atom can have required to remove the most loosely bound the same set of quantum numbers electron from an isolated gaseous form to a Aufbau Principle +1 ion (endothermic oxidation process) - electrons are filled up from the lowest to - a lower ionization energy means that the highest energy level element can be easily oxidized (charge Hund's Rule of Multiplicity becomes more positive) - each orbital of a given subshell is occupied - second ionization energy is always greater by a single electron before pairing begins than the first, because it takes more energy to remove an electron from a positive ion compared to a neutral atom CHEMICAL PERIODICITY - noble gasses have the highest ionization energies since they are generally nonreactive (not likely to give up an Electron Configuration electron) Electron Affinity - amount of energy absorbed when an electron is added to an isolated atom to form a negative ion charge (endothermic reduction) Ionic Radii - size: cations < neutral atoms < anions Electronegativity - measure of relative tendency of an atom to This form of the periodic table shows the sublevel attract electrons to itself when chemically blocks. Pale green inset: A summary of the sublevel combined with another element filling order. - highest electronegativity: F > O > N > CI Periodic Properties of the Elements CHEMICAL BONDING Basic Concepts of Chemical Bonding Ionic Bonding - electron transfer between atoms with large differences in their tendencies to lose or gain electrons (metal to non-metal) Covalent Bonding - electron sharing between atoms with small difference in their tendencies to lose or gain increasing value: ionization energy electronegativity electrons (non-metals) decreasing value: metallic character atomic and ionic radii electron affinity Lewis Symbols and the Octet Rule Metallic Character Octet Rule - luster, malleability, ductility, good conductor - MOST atoms lose/gain (via ionic bonding) of heat and energy or share (via covalent bonding) electrons to - Low ionization energy, electron affinity, try to attain a filled outer level of 8 electronegativity electrons - reactivity with halogens and oxygen 5 | UP KEM 2324 © KMCJ  University of the Philippines Chemical Engineering Society, Inc. (UP KEM) Chem 16 APE Module Drawing Lewis Structures - have the same relative placement of atoms - place one dot at a time on the four sides but different locations of bonding and lone (top, right, bottom, left) of the element electron pairs symbol. - arise due to delocalization of electrons - keep adding dots, pairing the dots until - resonance structures don’t actually all are used up. represent what the molecule looks like; in - metal: total number of dots is the maximum real life, the structure is a resonance hybrid number of electrons an atom loses to form (“bonds" represented with dashed lines) a cation. - formal charge of atom = no. of valence e- - - non-metal: number of unpaired dots equals (no. of unshared valence e + no. of bonds) either the number of electrons an atom Electronegativity and Bond Polarity gains in becoming an anion or the number - When one atom is significantly more it shares in forming covalent bonds. electronegative than the other atom it - note that the Lewis structure of a shares a covalent bond, it tends to pull the polyatomic ion is shown in square brackets, shared electron closer towards it, giving the with its charge as a right superscript bond a partially negative charge near the outside the brackets. electronegative atom and a partially positive charge in the opposite direction, and the bond is said to be polar Exceptions to the Octet Rule Bond Order - number of electron pairs being shared by a Electron-deficient molecules (with B or Be) pair of bonded atoms - lone pair-bonding pair > bonding pair-bonding pair 6 | UP KEM 2324 © KMCJ  University of the Philippines Chemical Engineering Society, Inc. (UP KEM) Chem 16 APE Module - See the appendix for geometry Molecular Orbital Diagram Electron-group Geometry - Regions of electron density around the central atom Molecular-group Geometry - arrangement of atoms around the central atom Valence Bond Theory - A covalent bond forms when orbitals of two atoms overlap and the overlap region, which is between the nuclei, is occupied by a pair of electrons - The space formed by the overlapping Bond Order orbitals has a maximum capacity of two - BO = ½ [(# of e- in bonding) - (# of e.- in electrons that must have opposite spins antibonding MO)] - The greater the orbital overlap, the stronger - Bond order > 0 implies that the molecular (more stable) the bond species is stable relative to the separate - π bonds: side by side overlap of orbitals atoms - σ bonds: head to head overlap of orbitals - Bond order = 0, no net stability and no - Single bonds: all are σ bonds likelihood that the species will form - Double bonds: 1 σ bond and 1 π bond - The higher the bond order, the stronger the - Triple bond: 1 σ bond and 2 π bond bond Hybridization Filling Molecular Orbitals with Electrons - The valence atomic orbitals in the molecule - In order of increasing energy are different from those in the isolated - Each orbital has a maximum capacity of 2e- atoms; different orbital shapes can “mix” or with opposite spins become hybridized - MOs of equal energy are half-filled, with spins parallel, before any of them is completely filled Magnetic Properties Paramagnetic - Has unpaired electron spins Diamagnetic - All electron spins are paired STOICHIOMETRY OF CHEMICAL REACTIONS Avogadro’s number - (NA) = 6.022 x 1023 - One mole = 6.022 x 1023 entities (molecules, atoms, ions, particles) Formula weight Molecular Orbital Theory - sum of the atomic weights of the elements in the formula - formation of molecular orbitals Stoichiometric factor Bonding - factor relating the moles of the required - region of high electron density between the reactant to moles of the other reactant nuclei due to molecular wave functions - comes directly from the coefficients in the being additive; high probability that the balanced chemical equation electrons are between the nuclei Antibonding aA + bB ⇌ cC + dD - region of zero electron density between the 1. Calculating mass of B given mass of A: nuclei due to waves canceling each other; zero probability for the electrons to be ( 𝑚𝑎𝑠𝑠 𝑜𝑓 𝐴 𝐹𝑊 𝑜𝑓 𝐴 ) × ( ) × (𝐹𝑊 𝑜𝑓 𝐵) = 𝑏 𝑎 𝑚𝑎𝑠𝑠 𝑜𝑓 𝐵 between the nuclei 2. Calculating volume of B given mass of A: ( 𝑚𝑎𝑠𝑠 𝑜𝑓 𝐴 𝐹𝑊 𝑜𝑓 𝐴 ) × ( ) × (22. 4 𝐿/𝑚𝑜𝑙) = 𝑏 𝑎 𝑣𝑜𝑙𝑢𝑚𝑒 𝑜𝑓 𝐵 7 | UP KEM 2324 © KMCJ  University of the Philippines Chemical Engineering Society, Inc. (UP KEM) Chem 16 APE Module 3. Calculating volume of B given volume of A: Single Displacement (𝑣𝑜𝑙𝑢𝑚𝑒 𝑜𝑓 𝑎) × ( )= 𝑏 𝑎 𝑣𝑜𝑙𝑢𝑚𝑒 𝑜𝑓 𝐵 - reactions where an element in a compound is replaced by another element. Percent composition 𝑚𝑎𝑠𝑠 𝑜𝑓 𝑒𝑙𝑒𝑚𝑒𝑛𝑡 Fe + Pb(NO3)2 → Pb + Fe(NO3)2 %𝑐𝑜𝑚𝑝 = 𝑡𝑜𝑡𝑎𝑙 𝑚𝑎𝑠𝑠 𝑜𝑓 𝑐𝑜𝑚𝑝𝑜𝑢𝑛𝑑 × 100% Double Displacement Empirical Formula - reactions where an element from each of - smallest whole number ratio of the atoms two compounds exchange places. present in a compound CaBr2 + AgNO3 → AgBr + Ca(NO3)2 Given: % composition of the compound Combustion Reactions 1. Assume 100g of sample. Mass of - reactions where the products are always each atom in the compound can be carbon dioxide and water. calculated. 2C2H2 + 5O2 → 4CO2 + 2H2O 2. Get the moles of each atom by dividing the masses with their Aqueous Ionic Reactions respective molecular weights. 3. Obtain the smallest number ratio of Precipitation reactions the moles of the element by dividing - wherein aqueous solutions combine to form the number of moles of each atom a precipitate; These reactions occur with the smallest number of moles. according to the solubility rules. Molecular Formula AgNO3 (aq) + NaCl(aq) → AgCl(s) + NaNO3 (aq) - actual numbers of atoms of each element Acid-Base Reactions present in a molecule of the compound - Reactions wherein an acid and base Given: % composition of the compound and combine and neutralize each other to form its actual molecular weight a salt. 1. Get Empirical Formula HBr(acid) + KOH(base) → H2O + KBr(salt) 2. Get the molecular weight of the Oxidation-Reduction Reactions empirical formula. - Reactions wherein the oxidation number of 3. Solve for n. an atom, molecule, or ion (redox reactions) n = (actual molecular changes by means of gaining or losing weight)/(empirical formula weight) electrons. 4. Multiply n to empirical formula Cl2 + 2KI → I2 + 2KCl Limiting Reactant/Reagent - the reactant that will all be consumed in the reaction Finding the limiting reactant: Redox Reactions Given: mass of Reactant 1 and Reactant 2 1. Get the mass of the product by Oxidation-Reduction (Redox) Reaction using the mass of reactant 1 - net movement of electron from one reactant 2. Get the mass of the product by to another using the mass of reactant 2. LEORA 3. Compare the two masses (product), Loss of Electrons, Oxidation, Reducing Agent the lower value will be the GEROA theoretical mass of the product Gain of Electrons, Reduction, Oxidizing Agent while the limiting reagent will be the Oxidation number (O.N.) reactant that gave that lower value. - the charge the atom would have if electrons Percent Yield were transferred completely %yield=(actual yield)/(theoretical yield) x100% - Oxidation: decrease in O.N. Types of Chemical Reactions - Reduction: increase in O.N. Combination - reactions where elements or compounds are joined together to form another compound. 2Na + Cl2 → 2NaCl Decomposition - Reactions where a compound breaks into smaller compounds or into elements. CaCO3 → CaO + CO2 8 | UP KEM 2324 © KMCJ  University of the Philippines Chemical Engineering Society, Inc. (UP KEM) Chem 16 APE Module Balancing Redox Reactions Heat capacity (C) 1. Divide the reaction into two half-reactions - quantity of heat required to change the 2. Balance the atoms in each half-reaction temperature of a system by 1 ℃ 3. Balance the charges by adding e-to the left q=C∆T side of the reduction half-reaction and to Specific heat (c) the right side of the oxidation half-reaction - heat capacity per gram 4. If necessary, multiply one or both half q=mc∆T reactions so that cH2O=4.18 J⁄(g ℃) number e- gained = number e- lost Law of conservation of energy 5. Add the balanced half-reactions, include - total energy remains constant during the states of matter interactions 6. Check that the atoms and charges are qsystem + qsurroundings=0 balanced Heat of reaction (qrxn) Balancing Redox reactions in acidic solution - quantity of heat exchanged between a 1. Divide the reaction into half-reactions system and its surroundings during 2. Balance atoms other that O and H chemical reaction at constant temperature 3. Balance O atoms by adding H2O molecules Exothermic reaction 4. Balance H atoms by adding H+ ions - reaction that gives off heat from system to 5. Balance charge by adding electrons the surroundings 6. If necessary, multiply one or both half qrxn0 - follows the same steps in balancing Bomb calorimeter reactions in acidic solution - a device used for measuring quantities of - add 1 mole of OH- to both sides of the heat, especially with combustion reaction for every H+ present qrxn= -qca; qcal=Ccal ∆T THERMOCHEMISTRY Laws of Thermodynamics Thermochemistry - concerned with interrelation of heat and First law of thermodynamics work that accompany chemical reaction - internal energy is the sum of all energies Internal energy U=q+w - total molecular energy internal to a - Endothermic, heat absorbed q>0 substance - Exothermic, heat released q0 - region of the universe that is studied - Work done by the system w0 - enthalpy change that occurs in the Third law of thermodynamics formation of one mole of the substance in - the entropy of a pure perfect crystal at 0 K the standard state is zero 𝑜 Standard molar entropy (S°) ∆𝐻 𝑓 = 0 𝑓𝑜𝑟 𝑁𝑎(𝑠), 𝐻2 (𝑔), 𝑁2 (𝑔), 𝑂2 (𝑔), 𝐶𝑔𝑟𝑎𝑝ℎ𝑖𝑡𝑒, 𝑎𝑛𝑑 𝐵𝑟2 (𝑙) - absolute entropy of one mole of a 𝑜 𝑜 𝑜 substance in its standard state ∆𝐻 𝑟𝑥𝑛 = ∑ 𝑣𝑝𝐻 𝑓 (𝑝𝑟𝑜𝑑𝑢𝑐𝑡𝑠) − ∑ 𝑣𝑅𝐻 𝑓 (𝑟𝑒𝑎𝑐𝑡𝑎𝑛𝑡𝑠) 𝑜 𝑜 𝑜 ∆𝑆 = ∑ 𝑣𝑝𝑆 (𝑝𝑟𝑜𝑑𝑢𝑐𝑡𝑠) − ∑ 𝑣𝑅𝑆 (𝑟𝑒𝑎𝑐𝑡𝑎𝑛𝑡𝑠) Bond-dissociation energy - quantity of energy needed to break one Gibbs’ Energy and Spontaneity mole of covalent bonds in a gaseous ∆G=∆H-T∆S species (kJ/mol) 𝑜 𝑜 𝑜 - Bond breakage +∆H ∆𝐺 = ∑ 𝑣𝑝𝐺 𝑓 (𝑝𝑟𝑜𝑑𝑢𝑐𝑡𝑠) − ∑ 𝑣𝑅𝐺 𝑓 (𝑟𝑒𝑎𝑐𝑡𝑎𝑛𝑡𝑠) - Bond formation -∆H - If ∆G>0 spontaneous process Lattice energy - If ∆Gion-dipole>ion-induced dipole>H-bonding> 𝐹α 2 dipole-dipole>London dispersion (dipole-induced 𝑑 Where dipole> induced dipole-dipole) q are the ion charges D is the distance Molecules can always have more than one type of - Coulomb’s Law determines: IMFA when it interacts with similar molecules. Melting point and boiling point of ionic compounds Comparing properties: look at the overall net effect Solubility of ionic compounds of the existing IMFA for the molecule as it acts in the - Ionic substance containing multiple entire substance charged ions usually have higher melting/boiling points than single charged Liquid Properties ions - For series of ions of similar charge Closer approach of smaller ions Surface Tension results in stronger interionic - Energy required to increase the surface attractions and higher area by a unit amount (J/m2) melting/boiling properties - Measure of the unequal attractions that Dipole-Dipole Interaction occur at the surface of the liquid - Occurs only for molecular compounds - Molecules at the surface are attracted Covalently bonded molecules unevenly Without charge - Stronger the IMFA, the greater surface Hydrogen Bonding tension - Only appears in molecules with H atoms Water has high surface tension due attached to N,O,F (three most to multiple H-bonds electronegative atoms) - Surfactants - Should consist of: Chemicals that decrease surface Hydrogen-bond donor: molecule tension of water by congregating at with attached hydrogen the surface and disrupting H-bonds F-H > O-H> N-H Capillary Action Hydrogen-bond acceptor: molecule - Rising of a liquid through a narrow space with the electronegative atom against the pull of gravity attracting the hydrogen - Capillarity N>O>F - Results from a competition between London Forces or Dispersion Forces intermolecular forces within a liquid - Weakest intermolecular forces 12 | UP KEM 2324 © KMCJ  University of the Philippines Chemical Engineering Society, Inc. (UP KEM) Chem 16 APE Module (Cohesive force) and those between the liquid and the tube walls (Adhesive force) Types of Solids According to Bond between - Capillary Rise: Adhesive force> Cohesive Particles Force (Water) - Capillary Fall: Cohesive Force > Adhesive Molecular Solids Force (Mercury) - Have molecules in each of the position of Vapor Pressure the unit cells - Pressure exerted by a liquid’s vapor on its - Held only by IMFAs surface at equilibrium - Fairly soft - Evaporation - Low melting point Process which molecules escape the - Volatile surface of a liquid and become gas - Electrical Insulator Temperature dependent - Ex. Ice, Sugar, Dry Ice, Naphthalene - Boiling point Network Covalent Solids Temperature at which liquid’s vapor - Have atoms that are covalently bonded to pressure is equal to the applied one another pressure - Very Hard - Normal Boiling Point - High Melting Point Boiling point at when the pressure - Poor thermal and electrical conductor is exactly 1 atm - Ex. Diamond, graphite, Sand (SiO2), SiC - High IMFA strength, Low Vapor Pressure Ionic Solids Viscosity - Have ions that occupy positions in a unit - Resistance to flow cell - Results from intermolecular attractions - Hard and brittle - High IMFA strength, High Viscosity - High melting point - Good thermal and electrical conductor Solid Properties when MOLTEN - Ex. CsCl, NaCl, ZnS Metallic Solids Dependence of properties on Structure: - Positively charged nuclei surrounded by sea - Rigidity/Hardness: Crystal Form/ Lattice of electrons - Melting Point: Strength of Atomic/Molecular - Soft to hard Interaction - Low to very high melting point - Conductivity: Electron Distribution along - Excellent thermal and electrical conductor entire solid - Malleable and ductile Amorphous Solid - Do not have well-ordered molecular Packing of Atoms structure - 1926 – Eugene Goldschmidt Crystalline Solid - Proposed atoms could be considered as - Well-defined structure that consist of packing in metallic solid as hard spheres extended array of repeating units - ABABAB Structure - Unit cells Hexagonal Closest Packed Each atom has 12 nearest neighbors Structure of Crystals 74.04% maximum filled space - ABCABC Structure Unit Cell Face Centered Cubic (Cubic closed - Smallest repeating unit of a crystal packing) Cubic Systems: Each atom has 12 nearest neighbor Simple Cubic Unit: 74.04% filled space - Each atom, ion, or molecule at a corner is Normal Melting Point shared by 8 unit cells. - temperature which solid melts at exactly 1.00 - 8x(1/8) = 1 particle atm pressure Body Centered Cubic - Higher IMFA strength, Higher melting point - Has an additional atom, ion, or molecule in Band Theory of Solids (Conductivity of Solids) the center of the unit cell - Extension of MO theory to solids - 8X(1/8) + 1 = 2 particles Face Centered Cubic - Unit cell has a cubic unit cell structure with an extra atom, ion, or molecule in each face - 8x (1/8) + 6x(1/2) = 4 particles 13 | UP KEM 2324 © KMCJ  University of the Philippines Chemical Engineering Society, Inc. (UP KEM) Chem 16 APE Module = 8.314 J/molK = 0.0821 L atm/molK Solving Gas Law Problem - Change in one of the four variables causes a change in the other while other two are held constant Ideal gas law reduces to one of the individual gas laws Units must be consistent: T: Kelvin - One variable is unknown and three are known, no change occurs - Insulators Ideal gas is directly applied to find Large gap between s and p bands the unknown Forbidden zone 1. Summarize the information: Identify - Semiconductors the changing gas variables- knowns Small gap between the bands and unknowns – and those held - Conductors or metals constant Have overlapping bands or partially 2. Predict the direction of change filled bands 3. Perform and necessary unit conversion 4. Rearrange the ideal gas law to GASSES obtain the appropriate relationship of gas variables and solve for the unknown variable Gas Laws Dalton’s Law of Partial Pressure - The pressure exerted by a mixture of gases Boyle’s Law is the sum of the partial pressures of the - Tackles about the relationship of volume individual gases and pressure Ptotal = Pa+Pb+Pc+….+Pn - Volume is inversely proportional to Pressure Law of Combining Volumes P1V1=P2V2 - When gases react, the volume of the Charles’ Law reactants and products, when measured at - about the relationship of volume and equal temperature and pressure, are always temperature whole number in ratios - Volume is directly proportional to Temperature Kinetic Molecular Theory 𝑉1 𝑉2 𝑇1 = 𝑇2 Combined Gas Law Equation Kinetic Molecular Theory of Gases - Combined Boyle’s and Charle’s Law in one - Based on three postulates: particle volume, statement motion, and collisions 𝑃1𝑉1 𝑃2𝑉2 Particle Volume 𝑇1 = 𝑇2 - Gas consists of a large collection of Gay-Lussac’s Law individual particles - about the relationship of pressure and - Volume of individual particle is extremely temperature small compared with the volume of its - Volume is directly proportional to container Temperature - Model pictures gas particles as point of 𝑃1 𝑃2 mass with empty space 𝑇1 = 𝑇2 - Proof: Gasses are easily compressible Avogadro’s Law Particle Motion - Volume is directly proportional to the moles - Gas particles are in constant, random, of gas straight-line motion, except when they - Standard molar volume collide with the container walls or with each Volume of 1 mol of gas at STP other 22.4L - Proof: Brownian Motion, molecular motion Ideal Gas Equation Particle Collisions - Combination of the other gas laws - Collisions are elastic PV = nRT - Colliding molecules exchange energy but Where they do not lose any energy through friction R can be derived at STP - Total kinetic energy is constant 14 | UP KEM 2324 © KMCJ  University of the Philippines Chemical Engineering Society, Inc. (UP KEM) Chem 16 APE Module - Proof: Sealed confined gas exhibits no (𝑃 + )(𝑉 − 𝑛𝑏) = 𝑛𝑅𝑇 2 𝑛𝑎 pressure drop over time 𝑉 2 Postulate 4 Where - KE of gas molecule is directly proportional a and b are correction constants to the absolute temperature a = intermolecular forces - Average KE of different gas molecules are b = volume of gas molecules equal at given temperature - Compressibility Factor - Proof: Brownian motion increases as 𝑍 = 𝑃𝑉 𝑛𝑅𝑇 temperature increases Ideal: Z=1 Average Translational Kinetic Energy Molecular volume is significant: Z>1 - Translational kinetic energy is directly Intermolecular forces of attraction: proportional to temperature only Z 0 ) M must be in kg/mol Solvation Graham’s Law - solvent-solute rxns - Diffusion - hydration if solvent is water Intermingling of gases ∆𝐻𝑠𝑜𝑙𝑢𝑡𝑖𝑜𝑛 = ∆𝐻𝑠𝑜𝑙𝑢𝑡𝑒−𝑠𝑜𝑙𝑢𝑡𝑒 + ∆𝐻𝑠𝑜𝑙𝑣𝑒𝑛𝑡−𝑠𝑜𝑙𝑣𝑒𝑛𝑡 + ∆𝐻𝑠𝑜𝑙𝑢𝑡𝑒−𝑠𝑜𝑙𝑣𝑒𝑛𝑡 - Effusion Crystal Lattice Energy Escape of gases through tiny holes - energy released when 1 mol of solid is - Rate of effusion is inversely proportional to formed from constituent ions in the gas the square roots of the molecular weights phase; measure of attractive forces in a or densities solid + − + − 𝑅1 𝑀2 𝑀 + 𝑋 → 𝑀 𝑋 + 𝑐𝑟𝑦𝑠𝑡𝑎𝑙 𝑙𝑎𝑡𝑡𝑖𝑐𝑒 𝐸 𝑅2 = 𝑀1 (𝑔) (𝑔) (𝑠) Miscibility 𝑅1 𝐷2 - dissolution of liquids in liquids 𝑅2 = 𝐷1 - Polar-polar and nonpolar-nonpolar liquids Real Gas Behavior are miscible - Polar gases are more soluble in water than Real Gas Behavior nonpolar gases (like dissolves like); polar - At low temperature and high pressure gases can form H bond in water and gases do not behave ideally (real gases) enhance their solubility - Reasons of deviation - Gases dissolve in liquids in exothermic Molecules are very close to one processes another thus volume is important Saturated solution Molecular interaction also became - equilibrium of dissolved and undissolved important solutes, where forward rate of rxn is equal to - Factors that causes deviation reverse rate of rxn Attractive forces Supersaturated sol’n: higher Nonpolar gasses attractive force is concentration than saturated sol’n London dispersion of dissolved solutes Polar gasses attractive force are Dissolution dipole-dipole attraction or hydrogen - Endothermic dissolution is enhanced by bonds heating Molecular weight is high - Exothermic dissolution is enhanced by Presence of lone pairs in the cooling molecule Solubility Rules - Van der Waals’ equation for the behavior of - All common compounds of Group I and real gasses at low temperatures and high ammonium ions are soluble. pressures 15 | UP KEM 2324 © KMCJ  University of the Philippines Chemical Engineering Society, Inc. (UP KEM) Chem 16 APE Module - All nitrates, acetates, and chlorates are - Because VP is lowered, temp must be raised soluble. to make VP equal atmospheric temp - All binary compounds of the halogens ΔTb = Kbxm (other than F) with metals are soluble, Osmotic pressure except those of Ag, Hg(I), and Pb. (Pb halides - Solvent passes from dilute sol’n to are soluble in hot water.) concentrated sol’n at a faster rate than in - All sulfates are soluble, except those of opposite direction, i.e. establishing barium, strontium, calcium, lead, silver, and equilibrium mercury (I). The latter three are slightly π = M RT soluble. Where - Except for rule 1, carbonates, hydroxides, π as osmotic pressure in atm oxides, silicates, and phosphates are The van’t Hoff Factor insoluble. - The relationship between the actual - Sulfides are insoluble except for calcium, number of moles of solute added to form a barium, strontium, magnesium, sodium, solution and the apparent number as potassium, and ammonium. determined by colligative properties Henry’s Law 𝑖 = 𝑎𝑝𝑝𝑎𝑟𝑒𝑛𝑡 # 𝑜𝑓 𝑝𝑎𝑟𝑡𝑖𝑐𝑙𝑒𝑠 𝑖𝑛 𝑠𝑜𝑙𝑛 # 𝑜𝑓 𝑚𝑜𝑙𝑒𝑠 𝑜𝑓 𝑠𝑜𝑙𝑢𝑡𝑒 𝑑𝑖𝑠𝑠𝑜𝑙𝑣𝑒𝑑 - The solubility of a gas in a liquid is directly proportional to the pressure of that gas above the surface of the solution ACIDS AND BASES 𝑃𝑔𝑎𝑠 = 𝑘𝑀𝑔𝑎𝑠 Quantitative Ways of Expressing Concentration 1. Percent mass of solute Acid-Base Theories 𝑚𝑎𝑠𝑠 𝑜𝑓 𝑠𝑜𝑙𝑢𝑡𝑒 %𝑤/𝑤 = 𝑚𝑎𝑠𝑠 𝑜𝑓 𝑠𝑜𝑙𝑢𝑡𝑖𝑜𝑛 * 100% 2. ppm Arrhenius Theory 𝑚𝑎𝑠𝑠 𝑠𝑜𝑙𝑢𝑡𝑒 (𝑚𝑔) - Acid: increases concentration of H+ ions, 𝑝𝑝𝑚 = 𝐿 𝑠𝑜𝑙𝑢𝑡𝑖𝑜𝑛 which form the hydronium ion H3O+ 3. Mole fraction - Base: increases concentration of OH− ions 𝑋𝑎 = 𝑚𝑜𝑙𝑒𝑠 𝐴 - Neutralization results from H+ and OH− 𝑡𝑜𝑡𝑎𝑙 𝑚𝑜𝑙𝑒𝑠 ions 4. Molarity 𝑚𝑜𝑙 𝑠𝑜𝑙𝑢𝑡𝑒 - Theory applies well only for aqueous 𝑀 = 𝐿 𝑠𝑜𝑙𝑢𝑡𝑖𝑜𝑛 solutions 5. Molality Bronsted-Lowry Theory 𝑚 = 𝑚𝑜𝑙 𝑠𝑜𝑙𝑢𝑡𝑒 - Acid: proton donor 𝑘𝑔 𝑠𝑜𝑙𝑣𝑒𝑛𝑡 - Base: proton acceptor Solubility of Gases - Neutralization as transfer of proton from an - Higher temp, less soluble acid to base - Higher pressure, more soluble - Reactions don’t have to occur in aqueous solutions Colligative Properties - Conjugate acid-base pairing Example H2O + NH3 → OH- + NH4+ Colligative Properties of Non Electrolytes Acid: H2O; Conj. Base: OH- Vapor pressure lowering Base: NH3; Conj Acid: NH4+ - Addition of nonvolatile solute to a sol’n Lewis Theory lowers its vapor pressure - Acid: electron-pair acceptor - Raoult’s Law: - Base: electron-pair donor Psolvent = Xsolute x P0solvent Example - Lowering of vapor pressure: CO2 + H2O ↔ NH2CO3 ΔPsolvent = P0solvent - Psolvent = P0solvent- Xsolvent x P0solvent ΔPsolvent =(1-Xsolvent ) x P0solvent Xsolute=1-Xsolvent Psolvent = Xsolute x P0solvent Freezing point depression - Addition of nonvolatile solute to sol’n lowers Acid: CO3 fpt. below that of pure solvent Base: H2O ΔTf = Kfxm Autoionization of Water and the pH Scale Boiling point elevation - According to Bronsted-Lowry, water can act - Addition of nonvolatile solute to a sol’n as either acid or base raises its bpt. above that of pure solvent 16 | UP KEM 2324 © KMCJ  University of the Philippines Chemical Engineering Society, Inc. (UP KEM) Chem 16 APE Module Amphoterism - ability of a substance to act as either acid CHANGES IN THE NUCLEUS or base Radioactivity and Nuclear Stability pH and pOH Types of Radioactive Emission Identical to pH and pOH Calculations Alpha Particles 4 α, 2α, 𝑜𝑟 2𝐻𝑒 4 2+ Helium-4 Kw = 1.00 x10-14 at 25 ℃ nuclei pH = - log [H3O+] pOH = - log [OH-] −1 0 pKw = - log [Kw]= 14 β , −1 β, High-speed Beta Particles pH + pOH =14 + 0 electrons 𝑜𝑟 β , β - acidic : pH < 7; pOH > 7 1 - Basic: pH > 7; pOH < 7 - Neutral: pH=pOH=7 Very high Gamma 0 Particles 𝛾 or 0γ energy photons Strengths of Acids and Bases NOTE: The weaker the acid/base, the stronger its conjugate (and vice versa) Modes of Radioactive Decay Mode ΔA ΔZ ΔN Strengths of Binary Acids - increases with decreasing H-X bond ⍺-decay -4 -2 -2 Bond strength: HF >>HCl > HBr > HI Acid strength: HF NH3 Electronegative H-X e− capture 0 -1 +1 - ↑ EN, ↑ Acidity Ternary Acids 𝛾 emission 0 0 0 - For ternary acids with same central element, ↑ Oxidation state of central element or ↑ O *see attached table for process atoms, ↑ Acidity HClO < HClO2 < HClO3 < HClO4 Nuclide stability - For ternary acids with same number of O - Key factors atoms, ↑ EN of central atom, ↑ Acidity 1. N, Z, N/Z ratio H2SeO4 < H2SO4 2. Total mass of nuclei HBrO4 < HClO4 - The points form a narrow band of stability HBrO3 < HClO3 that gradually curves above the line for N = Presence of Electronegative Atoms Z (N/Z = 1). - ↑ EN atoms, ↑ Acidity - Very few stable nuclides exist with N/Z < 1; 1 3 CH3COOH < CH2FCOOH < the only two are 1𝐻 and 2𝐻𝑒 CHF2COOH - Many lighter nuclides with N = Z are stable, Leveling Effect of Water 4 12 16 20 - Acid leveling effect masks differences in such as 2𝐻𝑒 , 6 𝐶, 8 𝑂, and 10𝑁𝑒 ; the heaviest acid strength in water 40 of these is 20𝐶𝑎. Thus, for lighter nuclides, - Strongest acid: H3O+ - Strongest Base: OH- one neutron for each proton (N = Z) is - Strong acids/bases react to form weak enough to provide stability. acids/bases - The N/Z ratio of stable nuclides gradually - Predicting Reactions: increases as Z increases. Thus, for heavier SA + SB = Neutral Sol’n stable nuclides, N > Z (N/Z > 1), and N WA + WB = Neutral Sol’n increases faster than Z. If N/Z of a WA + SB = Basic sol’n nuclide is either too high (above the band) e.g. CH3COOH + KOH → KCH3COO + or not high enough (below the band), the H2O nuclide is unstable and undergoes one of SA + WB = Acidic sol’n the three modes of beta decay. e.g. HClO4 + NH3 → ClO4− + NH4+ - All nuclides with Z > 83 are unstable and undergo ⍺ decay. 17 | UP KEM 2324 © KMCJ  University of the Philippines Chemical Engineering Society, Inc. (UP KEM) Chem 16 APE Module Balancing Nuclear Reactions 𝑡𝑜𝑡𝑎𝑙 𝐴 𝑡𝑜𝑡𝑎𝑙 𝐴 𝑡𝑜𝑡𝑎𝑙 𝑍 𝑅𝑒𝑎𝑐𝑡𝑎𝑛𝑡𝑠 = 𝑡𝑜𝑡𝑎𝑙 𝑍 𝑃𝑟𝑜𝑑𝑢𝑐𝑡𝑠 Kinetics of Radioactive Deacy Kinetics of Radioactive Decay - Decay Rate 𝑐ℎ𝑎𝑛𝑔𝑒 𝑖𝑛 # 𝑜𝑓 𝑛𝑢𝑐𝑙𝑒𝑖 𝑑𝑒𝑐𝑎𝑦 𝑟𝑎𝑡𝑒 = − 𝑐ℎ𝑎𝑛𝑔𝑒 𝑖𝑛 𝑡𝑖𝑚𝑒 - Units of Radioactivity - SI Unit: Becquerel (Bq) = 1 disintegration (d) per second - Curie (Ci) = 3.10 × 1010 d/s - For a large collection of radioactive nuclei, the number decaying per unit time is proportional to the number present; related by a decay constant k - The larger the value of k, the higher the decay rate - Radioactive decay is a first-order process - Half-life ( t1/2 ) = the time it takes for half the nuclei present in a sample or half of the mass of the sample to decay t1/2 = (ln 2)/k - NOT dependent on number of nuclei - inversely proportional to the decay constant, k - Nuclear binding energy - Energy required to break 1 mol of nuclei into their individual nucleons - Commonly expressed in electron volts (eV) or mega-electron volts (MeV) 1 eV = 1.602 × 10−19J = 10−6MeV - In Joules/mol, given change in mass in kg/mol, ΔE = Δmc2 - You can relate the change in mass to its energy equivalent using the conversion factor: 1 amu = 931.5 × 106 eV = 931.5 MeV - Calculating the binding energy per nucleon 𝑏𝑖𝑛𝑑𝑖𝑛𝑔 𝑒𝑛𝑒𝑟𝑔𝑦 𝐵𝑖𝑛𝑑𝑖𝑛𝑔 𝐸 𝑝𝑒𝑟 𝑛𝑢𝑐𝑙𝑒𝑜𝑛 = # 𝑜𝑓 𝑛𝑢𝑐𝑙𝑒𝑜𝑛𝑠 18 | UP KEM 2324 © KMCJ University of the Philippines Chemical Engineering Society, Inc. (UP KEM) Chem 16 APE Module 19 | UP KEM 2324 © KMCJ University of the Philippines Chemical Engineering Society, Inc. (UP KEM) Chem 16 APE Module 20 | UP KEM 2324 © KMCJ University of the Philippines Chemical Engineering Society, Inc. (UP KEM) Chem 16 APE Module 21 | UP KEM 2324 © KMCJ University of the Philippines Chemical Engineering Society, Inc. (UP KEM) Chem 16 APE Module 22 | UP KEM 2324 © KMCJ