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OME General Chemistry Lecture 1: Introduction, Terms, Symbols Dr. Hartwig Pohl Office: Beyer-Bau 122e Email: [email protected] Phone: +49 351 463 42576 1 Course Schedule 2015 Week Dates Mon Intro, Units, Charges, Atom Thurs Bohr Model, Periodic Table 1 Oct 19/22 2 Oct 26/29 no class Ionic & Covalent Bonds 3 Nov 2/5 Hybridization Reactions & Equations 4 Nov 9/12 Reactions & Equations Review (Take Home Quiz) 5 Nov 16/19 Gas Laws Kinetics 6 Nov 23/26 Thermodynamics Electrochemistry Review + Examples 7 Nov 30/Dec 3 Organic Molecules 8 Dec 7/10 9 Dec 14/17 Organic Molecules Review & Questions 2 Course Summary 1. Introduction 1.1 Preface 1.2 What is Chemistry? 1.3 History of Chemistry 1.5 Units, constants, symbols 2. Atomic models and Periodic table 2.1 General terms 2.2 Size and mass of atoms 2.3 Subatomic particles 2.4 Distribution of elementary particles in the atom 2.5 Number of elementary particles in the atom 2.6 Composition of the atomic nucleus 2.7 Composition of the electron shell 2.8 Periodic table of elements 3 Course Summary (cont.) 3. Chemical bonds and chemical formulas 3.1 Properties of compositions with different types of bonds 3.2 Ionic bonds 3.3 Covalent bonds 3.4 Metal bond 3.5 Composition of solids 3.6 Chemical formulas 3.7 Amount of substance and stoichiometry 4. Reactions and equations 4.1 Driving force of reactions 4.2 Reaction equations 4.3 Stoichiometry 4.4 Reactions of inorganic chemistry 4 Course Summary (cont.) 5. Gases 5.1 Gas law 5.2 Partial pressure 5.3 Real gases 5.4 Condensation of gases 6. Kinetics and Reaction mechanism 6.1 Kinetics and stability of chemical compounds 6.2 Definition of reaction rate 6.3 Reaction: 2. order 6.4 Reaction: 1. order 6.5 Reaction mechanism 6.6 Temperature dependence of reaction rate constant 6.7 Catalysis 5 Course Summary (cont.) 7. Thermodynamics 7.1 Equilibrium and mass action law 7.1.1 Temperature dependence of the equilibrium 7.1.2 Principle of the smallest constraint 7.2 Measurement of heat 7.3 Reaction enthalpy and entropy/ driving force of reactions 8. Electrochemistry 8.1 Conductivity and electrolysis 8.2 EMF 9. Organic Molecules Organic Chemistry (Dec 2015-Feb 2016) 6 Course Materials Lecture Slides note taking, reference Text Books L. Pauling “General Chemistry” or “College Chemistry” J.A. Campbell “Chemistry, the Unending Frontier” C.E. Mortimer “Chemistry” 7 What is Chemistry? Chemistry branch of science concerned with the substances of which matter (all things around us) is composed, the investigation of their properties and reactions, and the use of such reactions to form new substances. Properties characteristic qualities that can be observed without performing a chemical reaction Reaction process that converts a substances into other substances 8 Fields of Chemistry Analytical Chemistry Inorganic Chemistry Organic Chemistry Physical Chemistry Technical Chemistry Theoretical Chemistry Biochemistry Atmospheric Chemistry Nuclear Chemistry Food Chemistry Pharmaceutical Chemistry Photochemistry Polymer Chemistry Radiation Chemistry Macromol. Chemistry 9 History of Chemistry Prehistoric fire, pottery, tanning leather, dying cloth Ancient first definition of the atom as the simplest unit of matter Elements: fire, air, water, earth; properties: hot, cold, dry, wet 3500 BC glass, perfume, beer, wine, copper (from Malachite) in Egypt 2500 BC tin, iron, bronze (mixture of copper/tin) 1500 BC term element used in China 10 History of Chemistry 1500 BC dyeing of cloths Indigo (India) Alizarin (Egypt, Crete) 11 History of Chemistry Alchemy up to ca. 1500 AD driven largely by the desire to turn cheap metals into gold and to discover the Elixir of Life established a base for chemistry and periodic table 12 Fathers of Modern Chemistry 1600-1800 Dalton (Manchester) composition of chemical cmpds, term atom Berzelius (Austria, Stockholm) terms element, substance Boyle (London) “The Sceptical Chymist” (1661) collisions of particles in motion, more than 4 elements, encouraged experimentation Lavoisier (Paris) conservation of mass, chemical nomenclature, first chemistry textbook 13 First Synthesis of Organic Cmpds 1824 Oxalic acid by Wöhler, Göttingen uses: cleaning agent, extracting/isolating metals Friedrich Wöhler 1826 Aniline from Indigo Aniline is used in many modern dye synthesis and enabled the initial development of the dye industry. BASF = Badische Anilin- und SodaFabrik 1828 Urea by Wöhler AgNCO + NH4Cl → (NH2)2CO + AgCl first synthesis of a organic cmpd from inorganic precursor, believed only possible with living organisms 1840 Chemical fertilizer 14 Modern Chemistry 1865 Kekulé (Bonn) formula for benzene 1880 v. Baeyer (Leverkusen) synthesis of indigo 1870 L. Meyer (Tübingen), Mendelejew (Petersburg): Periodic system 1896 Becquerel (Paris) Radioactivity 15 Modern Chemistry 1905 Planck, Einstein (Berlin, Zürich): quantum theory 1906 1910 1912 Emil Fischer: studies on proteins Haber-Bosch method for ammonia production v. Laue (Berlin), Bragg (London), Debye, Scherrer (Berlin): use of X-rays for structural analysis 16 Modern Chemistry 1913 Bohr (Kopenhagen), Rutherford (Manchester): Atomic model 1939 1953 1960 Hahn, Strassmann (Berlin): nuclear fission (not really chemistry) Eigen (Göttingen): studies on very fast reactions t < 10–9 s Woodward, Fischer: Synthesis of complex natural materials, including chlorophyll, hemin, strychnine Calvin: mechanism of photosynthesis 1961 ... 17 Further Advancements • improvement and simplification of spectroscopic methods for analysis: UV, IR, NMR, X-ray • important findings in biochemistry: biochemical mechanisms • synthesis of peptides, crystallization and X-ray structures of proteins • functional polymers → modern plastics • surface analysis: scanning electron microscopy (SEM), atomic force microscopy (AFM) → resolution down to single atoms → surface analysis → catalysis; nanotechnology, Fullerenes • environmental protection, solar energy 18 Why Study Chemistry in OME? “It is through chemistry and her sister sciences that the power of man, of mind, over matter is obtained.” - Pauling Organic and Molecular Electronics (OME) Perspective Understand… • basic structure-property relationships and how they relate to functions in the device • how to tune device properties based on molecular structure and engineering • communicate effectively with colleagues 19 Matter & Radiant Energy The universe is composed of matter and radiant energy Matter all materials around us – gases, liquids, solids Radiant Energy light, x-rays, radio waves; ex: color of an object related to its absorption of light Chemistry is primarily concerned with the properties or characteristic qualities of matter rather than the particular object itself. 20 Types of Matter… Getting Specific Objects Heterogeneous Materials mixture… consists of parts with different properties; ex: granite Homogeneous Materials same properties throughout ex: quartz crystals Pure Substances Solutions homogeneous material with definite composition homogeneous mixture ex: sugar in water Elementary Substances substance that cannot be decomposed Compounds substance that can be decomposed into two or more other substances 21 Properties of Substances Properties characteristic qualities Physical Properties properties of a substance that can be observed without changing the substance into other substances ex: taste, solubility, density, melting point, malleability, ductility, hardness, color Chemical Properties properties that relate to its participation in chemical reactions 22 Energy Terms in Chemistry Energy involved in doing work or in heating an object; potential energy, kinetic energy, heat, radiant energy Chemical Energy energy from a chemical reaction which can do work; ex: mixture of gasoline vapor and air ignites, energy is liberated that can do the work of propelling a car, causes an increase in temperature of the engine and the exhaust gases Law of the Conservation of Energy whenever energy of one form disappears, an equivalent amount of energy in another form is produced Heat of the Reaction difference in heat contents of the products and the reactants 23 Units & Symbols SI-System (SI = Système international d´unités) Length Current Mass Temperature Time Amount m A kg K... (ºC) s mol Derived units, e.g. Force N = kg m/s2 Specific to Chemistry… Mass g Amount mol Molecular Weight g / mol Volume mL Solubility g / mL Density g / cm3 = g / mL Melting point ºC 24 Examples: Density A piece of gold (Mass = 301g) has a volume of 15.6 cm3. Calculate the density of gold. The density of mercury, the only liquid metal at room temperature, equals 13.6 g/cm3. Determine the volume of 5.50 g mercury. density = mass / volume density = mass / volume density = 301 g / 15.6 cm3 volume = mass / density density = 319.3 g / cm3 volume = 5.50 g / 13.6 (g/cm3) volume = 0.404 cm3 25 Example: Unit Conversion The density of gold is 19,300 kg m-3. Calculate the density in g cm-3. Density = Density = Density = Density = Density = 19.300 kg 1 m3 19.300 kg 1 m3 19.300 kg 1000 g 1 m3 1 kg 106 cm3 1000 g 1 m3 1 m3 1 kg 106 cm3 19.300 103 g 1 1 1 106 cm3 19.300 1g 1 103 cm3 = 19,3 g cm-3 26 1.5 values of basic constants, units and symbols Unit of atomic mass* u = L-1 g/mol = NA-1 g/mol = 1,6605655 × 10-27 kg Avogadro constant* NA = 6,022045 × 1023 mol-1 (= Loschmidt-number L) Bohr-Magneton B = eh/(4me) = 9,274078 × 10-24 J/Tesla Boltzmann constant* kB = 1,380662 × 10-23 J/K Electric constant 0 = 8,85418782 × 10-12 A2s2/(Jm) Electron mass me = 9,109534 × 10-31 kg Elementary charge* e = 1,6021892 × 10-19 As = 1,6021892 × 10-19 C Apparent gravity g = 9,81 m/s2 (average value) Faraday constant* FF = L × e = 96484,56 As/mol Gas constant* R = L × kB = 8,314472 J/(mol K) 27 Gravity constant G = 6,6720 × 10-11 m3/(kgs2) Core Magneton K = eh/(4mp) = 5,050824 × 10-27 J/Tesla Speed of light* c = 2,99792458 × 108 m/s (in vacuum) Magnetic constant 0 = 4 × 10-7 Vs/(Am) Planck constant* h = 6,626176 × 10-34 Js Proton mass mp = 1,6726485 × 10-27 kg * Particularly important for general chemistry Not all of these constants are needed for this lecture. Usually, calculation to 4-5 decimals is sufficient. 28 Letter capital_lower case A a B b G g D d E Z z H h Θ θ I i K k Λ λ Μ μ Ν v X x O o P R R Σ σ,ς* T τ Υ υ Φ n Χ χ Ψ ψ Ω ω *as final sound Name Alpha Beta Gamma Delta Epsilon Zeta Eta Theta Iota Kappa Lambda My Ny Xi Omikron Pi Rho Sigma Tau Ypsilon Phi Chi Psi Omega Pronounciation old greek new greek a a b w g g d th (soft) e e ds ds ä i t (th) th (hard) i i k k l l m m n n ks ks o o p p r r s s t t ü i f f ch ch ps ps oh oh 29 OME General Chemistry Lecture 2: Atomic Models & Periodic Table Dr. Hartwig Pohl Office: Beyer-Bau 122e Email: [email protected] Phone: +49 351 463 42576 30 The Atomic Theory John Dalton (1805, Manchester) All substances consist of small particles of matter, of several different kinds, corresponding to the different elements Atom Molecule smallest indivisible unit of an element atomos (Greek) = indivisible a group of atoms bonded to one another Arguments in Support of the Theory: • • • Democritus (460-370 BC) shared similar definition of atom with much less scientific backing Lavoisier and the demonstration of conservation of mass during a reaction Constant proportions: different samples of a substance contain its elements in the same proportions 31 Size & Mass of Atoms Modern microscopy methods enable to view atomic dimensions (scanning tunneling- or atomic force microscopy (STM or AFM)). Silicon Graphite Of course, it has been tried to further divide those particles, e.g. by collision with other particles or by irradiation with high energies. In all cases, it has been found that further division changes the properties. 32 Size & Mass of Atoms Pictures taken by scanning tunneling microscopy Nickel atoms on a copper surface A ring of cobalt atoms on a copper surface The size of atoms can in principle be taken from the pictures above. However, there are methods that are more precise and more simple, e.g. X-ray diffraction –→ Distances are in the order of 100–200 pm in case of atoms. 33 Size & Mass of Atoms Molecules are accordingly larger, usually a few hundred to a few thousand pm, some extremely large organic molecules may be even bigger. Mass of atoms and molecules Weighing is not possible due to small overall mass, but the methods described above help; e.g. 1 cm3 of an element X-ray diffraction → distances of atoms → number of atoms in 1 cm3, → mass of this number of atoms (Density ) → mass of an atom (Caution: different lattice types) Ex: Aluminum 4,489 × 10–23 g, Gold 3,27 × 10–22 g. Molecules can be accordingly heavier, but calculation with such small numbers is uncomfortable... atomic mass units u: 1 u = 1,6605655 × 10–24 g Ex: In this unit, the mass of an Aluminum atom is 26.89 u. 34 Subatomic Particles - Electrons Discovery of Electrons (Stoney) substances can be decomposed by an electric current, (Faraday) definite amount of electricity is needed to liberate a certain amount of element from a compound electricity exists in discrete units Properties of Electrons particle with negative electric charge (units of electric charge = Coulomb) magnitude of charge, q = -1.602 x 10-19 C mass, me = 9.107 x 10-28 g 1/10,000 as large as an atom! Flow of Electricity in a Metal amount (coulomb) rate (ampere = coulomb / second) rate of flow dependent on… potential difference / voltage drop btw ends 35 Subatomic Particles - Electrons 36 Determining the Charge of an Electron How big is the charge of an electron? – Millikan 1906 Oil Drop Experiment Spraying of oil droplets velocity in presence of gravity noted, then charged droplets were studied in the presence of a known electric field q = –1,602 × 10–19 C (Coulomb) = –1,602 × 10–19 As 37 Atomic Nuclei Identification of a Nucleus (Rutherford) every atom contains, in addition to one or more electrons, another particle called the nucleus of the atom every nucleus has a positive electric charge very small, very heavy (compared to an electron) nuclei are different for each element Protons simplest atomic nucleus charge exactly equal and opposite to that of an electron (+1.601 x 10-19 C) mass = mp = 1.672 x 10-24 g = 1836 x me = ca. 1 atomic mass unit Neutrons discovered by James Chadwick in 1932 mass = 1.675 x 10-24 g no electric charge! Nuclei are made up of protons and neutrons 38 Distribution of Particles in the Atom If the foil is a few µm thick, most a-particles penetrate through the foil without scattering. Evaluation: electrons must be distributed around the outer portion of the atom. Otherwise, far more particles must be scattered. 39 Distribution of Particles in the Atom a) Expected result b) Actual result Rutherford deduced from that (and from other experiments) that an electron will not diffract an a-particle from its track: → Matter consists mainly of empty space in which the point-shaped electrons move. The atomic nucleus including protons and neutrons contains almost the whole mass.: 1/1836 and is extremely small. 40 Approximate Size of an Atom The relative size of the nucleus with respect to the atom is approximately the same as a pea in the middle of this stadium. 41 Number of Particles in the Atom Larger amounts of matter must be neutral… NUMBER OF PROTONS (PZ) = NUMBER OF ELECTRONS (EZ) Each element has a particular number of protons. This is a definition of the term element on an atomic basis: Elements consist of atoms with an equal number of protons. The chemistry of an element is a function of its number of electrons/protons… Atomic number of the element (Z) = number of protons (= number of electrons) Further classification by mass number (MZ) = number of neutrons + protons number of protons = atomic number number of electrons = atomic number (in case of neutral particles) number of neutrons = mass number – atomic number (MZ – Z) 42 Number of Particles in the Atom Number of protons Number of electrons Number of neutrons Atomic number Mass number PZ = EZ EZ = PZ NZ = MZ – EZ Z = PZ = EZ MZ = PZ + NZ Element Name Symbol MZ Z EZ PZ NZ (MZ-Z) Sodium Na 23 11 11 11 12 Aluminum Al 27 13 13 13 14 Gold Au 197 79 79 79 118 43 Number of Neutrons: Isotopes Most elements appear with a variety of masses, i.e. the atoms contain the same number of protons but different numbers of neutrons. By suitable techniques, the atoms with different neutron numbers can be separated. These different atom types are called Isotopes. Isotopes are therefore elements with same numbers of protons but different numbers of neutrons. Most elements consist of a number of isotopes, sometimes mainly of one isotope, some of complex mixtures. e.g. hydrogen 99,99% MZ = 1 / 0,01 % MZ = 2 (Isotope with MZ = 2 has a name of its own = Deuterium) Chlorine Tin 75,5 % MZ = 35 / 24,5 % MZ = 37 Mixture of 10 Isotopes Isotopes of an element do not differ with respect to their chemical properties. 44 Isotopes of Sodium Two isotopes of Sodium: 23 11 Na und 24 11 Na 45 Information from Periodic Table 46