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Chapter 2 – Atoms and Elements What is chemistry? “A branch of science which deals with the elementary substances or forms of matter, of which all bodies are composed, the laws that regulate the combination of these elements in the formation of compound bodies, and the phenomena that accompany their exposure to diverse physical conditions”. Composition Preparation Reaction What is it made of ? How is it made? How does it interact with others or react with its surroundings ? Ex) COFFEE Composition: i) Organic Compounds: Proteins Esters Acids Sugars Caffeine Pesticides ii) Inorganic Compounds: Water Dissolved Salts Dissolved minerals 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 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. 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. Where does chemistry fit? What is the purpose of modern chemistry? 1) Physical chemistry/chemical physics Thermodynamics Kinetics Spectroscopy Quantum Mechanics 2) Analytical/Environmental Chemistry Quantification and identification techniques Separation Methods Development of New instrumentation Forensic/Environmental Chemistry 3) Preparative Chemistry Synthesis of new organic compounds Synthesis of new inorganic compounds Material Science Pharmaceutical Chemistry 4) Biochemistry/Molecular Biology Chemical processes of Life Structure and function of proteins Exploration of DNA Rational Drug Design Atomic Theory Greeks Atom ( A – not, tomos – to cut) Plato Aristotle - Revelation of truth through logic - Cosmic order - Hierarchy of being Atomic Theory Greeks Five perfect shapes Tetrahedron Cube Octahedron Dodecahedron Icosahedron Five elements Fire Water Wind Earth Ether Technology Steam Engines Organs Jewelry Reinforced Concrete Medieval Times Religious Society Communal Static God Centered , Hierarchical No Individual Identity Cycles of Life and Nature Preordained unchangeable Order No Change “Understanding God’s Creation” “Moral Teachings From Nature” Alchemy Transmutation of lead into gold. Chemical knowledge of the time Coal/Peat Fuel Salt Production Pigments/Dyes Poisons/Medicines Metal work: - steel - pewter - quicksilver - jewelry Beer/Wine Distillation Fragrances/Extracts Glass/Ceramics Enlightenment Scientific Method Determinism Materialism Mechanistic Thinking Earth Centered Individualism Career Scientist “ Mechanistic Understanding of the Universe” Times of Change/Discovery - French and American Revolution. - Industrial Revolution - Rapid exploration of chemistry began: New Elements Natural Products Synthetic Methods Lavoisier Joseph Proust John Dalton 1785 “Conservation of mass” 1794 “Law of Definite Proportions” 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 identical in 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. Modifications Required to Daltons Theory 1. Atoms can be further divided into 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. Elemental Forms Molecular Gas Molecular Liquid Ex) Bromine is a liquid composed of diatomic molecules Br2. Ex) Chlorine is a gas composed of diatomic molecules Cl2. H2, N2, O2, F2, and Cl2. Elemental Forms Atomic Gases Molecular Solids Ex) I2, P4, and S8. Atomic Liquids Ex) He, Ne, Ar, Kr, Xe, and Rn . Infinite Networks Ex) Metals Diamond Graphite Ex) Hg, Ga Graphite (2-D) Diamond (3-D) Compounds Molecular Infinite networks – SiO2 Sand, Glass Ionic Complex ions are also possible such as: SO32-, NO3-, PO4- NH4+, H3O+ Properties Chemical Is observed by changing a compound/element into another compound/element. Physical Is observed without changing a compound/element into another compound/element Chemical Reactions Melting point Energy of Reaction Freezing Point Combustion Density 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 Henri Becquerel 1896 Marie and Paul Currie 1899 Proton/Nucleus Ernest Rutherford 1919 Neutron J. Chadwick 1932 Radioactivity 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 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 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 “Canal Rays” and Protons e- e- e - + + + Cathode + Anode E Goldstein (1850-1930) discovered canal rays in 1886using 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 Electrons emitted from the cathode hit gas molecules causing ionization into (more) electrons and leaving positively charged “ions” which travel to the cathode E. Rutherford determined that the hydrogen cation is a fundamental particle, and named it the proton Radioactivity Three types of radiation: alpha, a , beta, b, and 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. Radioactivity: properties From their charge-mass ratios and other experiments of these rays were characterized and identified Alpha particles: He2+ nuclei m = 4 amu q =+2) Beta particles: electron (e-) (identical to cathode rays) Gamma rays: high-energy light, with wavelengths shorter than X-rays Rutherford experiment Using alpha particles, he bombarded a very thin foil of gold and observed deflections using a circular fluorescent screen 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. Size of 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. 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 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 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 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 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 equal to the combined number of protons and neutrons it contains. Mass number (A) = # protons + # neutrons mass number symbol atomic number If # p’s = #e’s neutral 12 6 C 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). Exercise e.g. Gallium has two naturally occurring isotopes and an average atomic mass of 69.723 u: 69 71 68.926 u 70.925 u G G 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 % The Mole It is not practical to work on the scale of individual atoms. It is necessary to work on the macroscopic scale. It was found that for 6.0221*1023 atoms for any element the mass corresponds to the atomic mass in grams. This number, 6.0221*1023, was named after Amedeo Avogadro who initialyl proposed the idea. Ex) 12.00 g of carbon corresponds to 6.02221*1023 carbon atoms The same is true for molecules. Ex) CO2 weighs 12.000 + 2*15.999 = 43.998 u 6.02221*1023 molecules of CO2 weighs 43.998 g “1 mole = 6.0221 × 1023” Molar Mass The molar mass of a particle is the mass in grams of one mole, 6.0221*1023, particles Ex ) One mole of protons weighs (6.0221*1023)*(1.67*10-24 g) = 1.01 g Ex) CO2 has molecular mass of 43.998 u therefore, it has a molar mass of 43.998 g/mol Exercise: How many moles of Cl2 are there in 105.7g. The atomic mass of Cl is 35.46 u The molecular mass of Cl2 is 2*35.46 u = 70.92 u Its molar mass is 70.92 g/mol. # moles Cl2 = mass Cl2/molar mass Cl2 =105.7g/70.92 g/mol = 1.490 moles Exercise: What is the mass of carbon dioxide containing 2.57*1021 atoms of oxygen CO2 contains two O atoms for every CO2 molecule # CO2 = (# of O)/2 = 2.57*1021/2 = 1.29*1021 # moles CO2 = (# CO2 molecules )/( 6.0221*1023 molecules/mol) # moles CO2 = (1.29*1021 molecules)/( 6.0221*1023 molecules/mol) # moles CO2 = 0.00214 mol How much does this weigh? mass CO2 = (# moles CO2)*(molar mass of CO2) mass CO2 = (0.00214 mol)*(43.998 g/mol) = 0.0916 g Law of Periodicity Group Similar chemical properties “The properties of the elements are periodic functions of atomic number.” Period Repetition of properties Nonmetals – insulators not ductile Metalloids - Semiconductors Ductile ? Metals – Conducting, Ductile Overview of the Elements by Group Hydrogen (H) Has properties of groups 1 and 17 but doesn’t belong to either. Diatomic gas (H2) Unreactive Group 1: Alkali Metals (Li, Na, K, Rb, Cs, Fr) Soft metals that react strongly with Water and oxygen. (reactivity increases with atomic mass) Do not exist in pure form in nature due to high reactivity Readily lose 1 electron to make cations with +1 charge Group 2: Alkaline Earth Metals (Be, Mg, Ca, Sr, Ba, Ra ) Most are metals that react with water to give X(OH)2 and with oxygen to give XO. Reactivity increases with atomic mass. Beryllium does not react with water (Highly toxic) Do not exist in pure form in nature due to high reactivity Readily lose 2 electrons to make cations with +2 charge 3 Groups 3-12: Transition Metals 4 5 6 7 8 9 10 11 12 Scandium Titanium Vanadium Chromium Manganese Iron Cobalt Nickel Copper Zinc 44.9559 47.88 50.9415 51.9961 58.9332 58.693 63.546 65.39 Sc Ti V 21 22 Yttrium Zirconium Niobium 88.9059 91.224 Y 39 23 Zr 55.847 Mn 25 Fe 26 Co 27 Molybdenum Technetium Ruthenium Rhodium 95.94 (98) 101.07 102.906 Mo Nb Tc Ru Ni Cu Zn 28 29 30 Palladium Silver Cadmium 106.42 107.868 112.411 Rh Pd Ag Cd 40 41 42 43 44 45 46 47 48 Hafnium Tantalum Tungsten Rhenium Osmium Iridium Platinum Gold Mercury 178.49 180.948 183.85 186.207 190.2 192.22 195.08 196.967 200.59 La-Lu Hf Ta W Re Os Ir Pt 72 73 74 75 76 77 Rutherfordium Dubnium Seaborgium (263) Bohrium Hassium Meitnerium Darmstadium (262) (265) (266) (261) Ac-Lr Cr 24 92.9064 54.9380 Rf 104 (262) Db 105 Sg 106 Bh 107 Hs 108 78 Mt 109 (281) Au Hg 79 80 (1994) (272) (1996) (285) 111 112 Dt 110 Group 3 metals lose 3 e- to make +3 cations otherwise act like Group 2 metals Groups 4-11 are the ‘true’ transition metals in that they lose e- to form coloured compounds in which the metal atom has a positive charge Groups 11 and 12 mimic the behaviour of Groups 1 and 2 respectively but are less reactive The metals in the middle of the transition groups are hardest Ex) group 6 Cr, Mo, W The metals at the edges are the softest (Groups 3 and 12) Ex) Zn and Ga Readily lose electrons to make cations Mercury is the only liquid metal; the rest are solids Copper and gold are the only coloured metals Gold is the most malleable metal Silver is the best conductor Group 13 (B, Al, Ga, In, Tl) Most are metals; Boron is a metalloid All are solids, Gallium -low m.pt. Aluminum - industrially important - third most abundant - produces its own protective layer Lose 3 electrons to make cations with +3 charge Form compounds in a 1:3 ratio with halogens (e.g. BCl3) Group 14 (C, Si, Ge, Sn, Pb) Form compounds in a 1:4 ratio with halogens (e.g. CCl4) C is a nonmetal; Si and Ge are metalloids; Sn and Pb are metals Si is the 2-nd most abundant Element which does occur in pure form but as silicates (compounds made of silicon and oxygen) which form rocks, sand, glass, etc. CARBON Carbon exists in several different Allotropes: 1) graphite 2) diamonds 3) fullerenes – Many types Carbon is the backbone atom of organic and biological molecules Group 15: Pnictogens (N, P, As, Sb, Bi) N and P are nonmetals As and Sb are metalloids Bi is somewhat metallic N is a highly stable diatomic gas (N2)and the most abundant element in the atmosphere P in three allotropes White, P4 – Fire Bombs Red and black, polymers - used in match heads) Form compounds in a 1:3 ratio with hydrogen (e.g. NH3) Group 16: Chalcogens (O, S, Se, Te, Po) O, S & Se are nonmetals Te is a metalloid Po is a metal O is the most abundant in the earth’s crust & the second most in the atm. O exists in two allotropes : O2 and O3 both are very reactive gases S exists in many allotropes: S2, S6, S8, etc. Form compounds in a 1:2 ratio with hydrogen (e.g. H2O) Gain 2 electrons to make anions with -2 charge Group 17: Halogens (F, Cl, Br, I, At) Nonmetals that exist as diatomic molecules (except for astatine which is too unstable to study) F & Cl are gases Br is a liquid Iodine is a solid colourful F2 is yellow Cl2 is yellow-green Br2 is red-brown I2 is dark purple l2 Br2 Cl2 F most reactive known Form compounds in a 1:1 ratio with hydrogen (e.g. HF) Gain 1 electron to make anions with -1 charge Group 18: Noble Gases (He, Ne, Ar, Kr, Xe, Rn) Inert- unreactive gaseous nonmetals Exist primarily in elemental from Compounds have been made containing Xe Ex) XeF2 Helium low density – Balloons Low BP - Coolant Xenon Used as a probe to study structure in porous material Glow when a current passes through them: Ex) Neon Lights Concepts Dalton’s atomic theory of matter Elemental Forms Chemical & physical properties Subatomic particles (protons, neutrons, electrons) Models of the atom Isotopes, calculating average atomic mass and percent abundance Atomic number and mass number Elements (names and symbols) Avogadro’s number and the mole Periodic table (groups and periods)