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What is chemistry?

What is chemistry?.

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What is chemistry?

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  1. What is chemistry? “A branch of science which deals with: i)theelementary substances or forms of matter, of which all bodies are composed; ii)the laws that regulate the combination of these elements in the formation of compound bodies; and iii)thephenomenathat accompany their exposure to diverse physical conditions”. Composition What is it made of ? Preparation How is it made? How does it interact with others or react with its surroundings ? Reaction

  2. COFFEE Composition: i) OrganicCompounds: ii) InorganicCompounds: Proteins Water Esters Dissolved Salts Acids Dissolved minerals Sugars Caffeine Pesticides

  3. COFFEE Preparation Grown • Biochemical processes make the organic & biological • compounds Roasted • Heat combined with air burns off undesired compounds & • converts some to those that give flavor • Caffeine is burned off if roasted too long • Decaffeination

  4. COFFEE Preparation Ground • Pulverization of the bean to increase the surface are to aid • extraction process • Makes it more vulnerable to oxidation affecting taste & shelf • life Extraction • Hot water poured over powder, where all water soluble • compounds dissolve. • The liquid is separated from the bean • residue by a filtration.

  5. COFFEE Reaction with Surroundings Caffeine • Stimulant – increases heart rate by promoting adrenaline • production • Diuretic – stimulates urine production Burns • When it sits the element exposed to the air the organic compounds oxidizes causing a bitter taste.

  6. Atomic Theory Greeks Atom ( A – not, tomos – to cut) Plato - Revelation of truth through logic • Cosmic order • Hierarchy of being Aristotle

  7. Atomic Theory Greeks Five perfect shapes Fiveelements Tetrahedron Cube Octahedron Dodecahedron Icosahedron Fire Water Wind Earth Technology Ether Steam Engines Organs Jewelry Reinforced Concrete

  8. Enlightenment ScientificMethod “ Mechanistic Understanding of the Universe” Determinism Mechanistic Thinking Materialism Earth Centered Individualism Career Scientist Times of Change/Discovery - French and American Revolution. - Industrial Revolution • Rapid exploration of chemistry began: • New Elements • Natural Products • Synthetic Methods

  9. Lavoisier 1785 “Conservation of mass” Joseph Proust 1794 “Law of Definite Proportions” John Dalton 1808 “Atomic Theory of Matter” 1. All matter consists of solid and indivisible atoms. 2. All of the atoms of a given chemical element are identicalin mass and in all other properties. 3. Different elements have different kinds of atoms; these atoms differ in mass from element to element. 4. Atoms are indestructible & retain their identity in all chemical reactions. 5. The formation of a compound from its elements occurs through the combination of atoms of unlike elements in small whole-number ratios.

  10. Modifications Required to Daltons Theory 1. Atoms can be further dividedinto subatomic particles. Ex) Protons, neutrons, electrons 2. Different isotopes of an element have different masses Ex) Carbon-12 12.000 u Carbon-13 13.003 u Carbon-14 14.003 u 3. Valid, However some have very similar masses. Ex) Nitrogen-14 14.003 u. Carbon-14 14.003 u. 4. In nuclear reactions, atoms do not retain their identity. Ex) Radium-226 → Radon-222 + a-4 5. Valid, however, Dalton was unaware that not all elements are made up of single atoms.

  11. Modern Atomic Theory In the late 19-th and early 20-th century the basic principles of modern atomic theory were laid down Electron J.J. Thomson 1896 R. A. Millikan 1909 Radioactivity Henri Becquerel 1996 Marie and Paul Currie 1899 Proton/Nucleus Ernest Rutherford 1919 Neutron J. Chadwick 1932

  12. Electrons Hole drilled in tube. Gass entering tube glows Cathode Ray Tube Cathode: negative electrode Anode: positive electrode Current flows when tube is evacuated Cathode Rays

  13. Electron charge-to-mass ratio J.J. Thomson – 1897 - cathode rays are negatively charged particles CRT with electric and magnetic fields applied at right angles Beam deflects to positively charged plate Magnetic field applied to deflected beam Changes in the deflection behaviour allowed the mass to charge ratio of the electron to be determined at 1.7588202 C/kg

  14. Oil Drop Experiment R Millikan and H A Fletcher (1909) Accurate measurement of the electron charge. Balanced the force of gravity with an opposing electric force The balancing force between droplets had common factor He surmised that the charge of a single electron e = 1.60217646 10-19 C Applying the charge/mass ratio,mass of e = 9.1093819 10-31 kg

  15. “Canal Rays” and Protons - Cathode + e- + e- + e- + Anode Electrons emitted from the cathode hit gas molecules causing ionization into (more) electrons and leaving positively charged “ions” which travel to the cathode - cations. E Goldstein (1850-1930) discovered canal rays in 1886 using a “reverse cathode ray” tube Those that pass through the hole (“canal”) can be analyzed for charge-mass ratio, which are much smaller than electron, but largest for hydrogen E. Rutherford determined that the hydrogen cation is a fundamental particle, and named it the proton

  16. Radioactivity Three types of radiation: alpha, a , beta, b, & gamma, g. Paul and Marie Currie isolated the radioactive elements Radium and Polonium. They postulated that their spontaneously emitted radiation was the result of nuclear disintegration. Three fundamental types of nuclear radiation were identified by how they respond to electric fields by E. Rutherford.

  17. Radioactivity: properties From their charge-mass ratios and other experiments of these rays were characterized and identified Alpha particles: He2+ nuclei m = 4 amuq =+2) Beta particles: electron (e-) (identical to cathode rays) Gamma rays: high-energy light, with wavelengths shorter than X-rays

  18. Rutherford experiment Using alpha particles, he bombarded a very thin foil of gold and observed deflections using a circular fluorescent screen

  19. The nuclear atom He tried to prove the plum pudding model of the atom propose by Thomson, which is composed of electrons imbedded in a sphere of uniform positive charge. Rutherford said of the alpha particles deflected almost straight back. • Deflection angle and frequency were carefully • measured, which led to the conclusions: • 1. Most of gold foil is empty space • 2. There are small centers of highly-positive charge • 3. Centers have high mass to resist displacement • 4. Sizeof atom estimated from distance between centers to be ~10-10 m diameter. • 5. Size of centers estimated to be ~10-15 m diameter • Centers were called the nucleus. • Electrons occupy the volume of the atom outside • the nucleus

  20. Constituents of the atom In 1920 Rutherford predicted the existence of the neutral particle with mass equal to that of a proton and electron. In 1932 Chadwick verified experimentally the existence of the neutron Relative mass of carbon defined t be 12 u

  21. The mass spectrometer Mass spectrometer is a variation on the CRT, developed by J.J. Thomson, which allows the determination of m/z ratios of cations. Cations of differing m/z ratio’s can be selected by adjusting the magnetic field strength

  22. Average atomic mass Isotopes are atoms of the same element that differ in mass due to differences in the number of neutrons 35Cl has 17 protons and 18 neutrons 37Cl has 17 protons and 20 neutrons Similar to computing an average grade of a class The atomic mass of Chlorine is a weighted average between the two isotopes as: Atomic Mass = Mass(Cl-35) *frac.(Cl-35) + Mass(Cl-37) *frac.(Cl-37) = (34.968)*(0.7537) + (36.956)*(0.2463) = 35.46 u

  23. Defining an Element The atomic mass unit (u)is defined as one twelfth of the mass of a carbon atom containing six protons, six neutrons and six electrons: 1 u = 1.661 × 10-24 g The mass of an atom in u will be approximately equalto the combined number of protons and neutrons it contains. Mass number (A) = # protons + # neutrons If # p’s = #e’sneutral If # p’s > # e’s cation If # p’s < # e’s anion Atomic number (Z) = # protons The atomic # determines the identity of the element (optional).

  24. Exercise Gallium has two naturally occurring isotopes and an average atomic mass of 69.723 u: 69G 71G Atomic masses measured using Mass Spectrometry 68.926 u 70.925 u Calculate the percent abundance of each isotope of gallium. At. Mass = M(69G)*frac(69G) + M(71G)*frac(71G) frac(69G) + frac(71G) =1 frac(69G) =1- frac(71G) =1-x At. Mass = M(69G)*(1-x) + M(71G)*x 69.723 = (68.926)*(1-x) + (70.925)*x= 68.926+1.999*x x =(69.723-68.926)/1.999 = 0.3987 = 39.87 %

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