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2 The Atomic Nature of Matter Introduction: The Evolution of Atomic Theory 2.1 - Atomic Theory 2.2 - Atomic Architecture: Electrons and Nuclei 2.3 – Atomic Diversity: The Elements 2.4 – Charged atoms: Ions 2.5 – Energy of Atoms and Molecules The Greek Concept of Atomos: The Indivisible Atom • Around 440 BC, Leucippus originated the atom concept. • His pupil, Democritus (c460-371 BC) extended it • There are five major points to their atomic idea. • All matter is composed of atoms, which are too small to be seen. These atoms CANNOT be further split into smaller portions. • There is a void, which is empty space between atoms. • Atoms are completely solid. • Atoms are homogeneous, with no internal structure. • Atoms can differ in size, shape, and weight Aristotle (384-322 BC) Spoke openly against the concept — Only a few scholars gave it much thought. The Catholic Church accepted Aristotle's position — the atom concept diminished equated atomistic ideas with Godlessness It was not until 1660 that Pierre Gassendi succeeded in separating the two not until 1803 that John Dalton put the atom on a solid scientific basis. John Dalton 1803-1807 Modern Atomic Theory 1) All matter is composed of tiny particles called atoms. 2) All atoms of a given element have identical chemical properties that are characteristic of that element. 3) Atoms form chemical compounds by combining in whole-number ratios. 4) Atoms can change how they are combined, but they are neither created nor destroyed in chemical reactions. 2.1 Atomic Theory Conservation of Atoms & Mass Atoms are neither created nor destroyed during physical or chemical processes. Fig 2-4 2.1 Atomic Theory Conservation of Atoms & Mass Mass is neither created nor destroyed during physical or chemical processes. Fig 2-5 Courtesy Patrick Watson 2.1 Atomic Theory Atoms & Molecules are Continually in Motion Diffusion: the gradual mixing of atoms & molecules due to their continual motion. Fig 2-7 2.1 Atomic Theory Atoms & Molecules are Continually in Motion Fig 2-9 Courtesy Patrick Watson 2.1 Atomic Theory Atoms Combine to Make Molecules Fig 2-6 2.1 Atomic Theory Dynamic Molecular Equilibrium the condition in which a forward and reverse process occur at equal rates, so the system undergoes no net change. Fig 2-10 When this example is at dynamic equilibrium, the number of molecules vaporizing equals the number of molecules condensing at any given moment. Laws of Mass Conservation & Definite Composition Law of Mass conservation: The total mass of substances does not change during a chemical reaction. Law of Definite ( or constant ) composition: No matter what its source, a particular chemical compound is composed of the same elements in the same parts (fractions) by mass. Mass Percent Composition of Na2SO4 Na2SO4 = 2 atoms of Sodium + 1 atom of Sulfur + 4 atoms of Oxygen Elemental masses Percent of each Element 2 x Na = 2 x 22.99 = 45.98 1 x S = 1 x 32.07 = 32.07 4 x O = 4 x 16.00 = 64.00 142.05 Check % Na = Mass Na / Total mass x 100% % Na = (45.98 / 142.05) x 100% =32.37% % S = Mass S / Total mass x 100% % S = (32.07 / 142.05) x 100% = 22.58% % O = Mass O / Total mass x 100% % O = (64.00 / 142.05) x 100% = 45.05% % Na + % S + % O = 100% 32.37% + 22.58% + 45.05% = 100.00% Calculating the Mass of an Element in a Compound Ammonium Nitrate How much Nitrogen is in 455 kg of Ammonium Nitrate? Ammonium Nitrate = NH4NO3 The Formula Mass of Cpd is: 4 x H = 4 x 1.008 = 4.032 g 2 x N = 2 X 14.01 = 28.02 g 3 x O = 3 x 16.00 = 48.00 g Therefore gm Nitrogen/ gm Cpd 80.052 g 28.02 g Nitrogen = 0.35002249 g N / g Cpd x 100 = 35.00% 80.052 g Cpd 455 kg x 1000g / kg = 455,000 g NH4NO3 455,000 g Cpd x 35.00 g N / 100g Cpd = 1.59 x 105 g Nitrogen or: 455 kg NH4NO3 X 28.02 kg Nitrogen = 159 kg Nitrogen 80.052 kg NH4NO4 Law of Multiple Proportions If elements A and B react to form two compounds, the different masses of B that combine with a fixed mass of A can be expressed as a ratio of small whole numbers: Example: Nitrogen Oxides I & II Nitrogen Oxide I : 46.68% Nitrogen and 53.32% Oxygen Nitrogen Oxide II : 30.45% Nitrogen and 69.55% Oxygen Cmpd I Cmpd II in 100 g of each Compound: g O = 53.32 g & 69.55 g g N = 46.68 g & 30.45 g g O /g N = 1.142 & 2.284 Cmpd II 2.284 2 = Cmpd I 1.142 1 2.2 Atomic Architecture: Electrons & Nuclei Forces Gravitational force: the force which pulls object toward the center of the Earth. Electrical force: the attraction or repulsion between two charged objects. Fig 2-12 2.2 Atomic Architecture: Electrons & Nuclei Forces Gravitational force: the force which pulls object toward the center of the Earth. Electrical force: the attraction or repulsion between two charged objects. Magnetic force: the force generated by charged objects in motion. Fig 2-12 Courtesy Patrick Watson Davy, Faraday 1807 Humphry Davy: • Forces holding compounds together are electrical 1833 Michael Faraday: — Atomic mass and the electricity needed to free elements from compounds are related More Important developments • 1864 Heinrich Geissler: developed a pump • produced good vacuums in sealed glass tubes • 1870’s William Crookes: Cathode rays • are negatively charged • same regardless of cathode metal • are particles with mass • 1891 George Stoney • electricity exists in units called "electrons” 2.2 Atomic Architecture: Electrons & Nuclei Electrons Fig 2-13 Gas Discharge Tube Conclusions: •Atoms are made up of smaller positive and negative fragments. •The negatively charged particles are electrons, which are uniform in behavior, regardless of their source. 1897 J.J. Thomson DISCOVERED THE ELECTRON 1897 J.J. Thomson • Deflected an electron beam by both magnetic and electric attraction/repulsion • Measured the electron’s mass/charge ratio 2.2 Atomic Architecture: Electrons & Nuclei Electrons Fig 2-14 Cathode Ray Tube Conclusion: Charge e 1.76x1011 C/kg Mass m Cathode Rays • • • • Attracted to the positive electrode Not visible but could make things “glow” Traveled in a straight line Could be bent by electric or magnetic fields • A plate in it’s path acquired a negative charge • Same regardless of material The quest continues • 1905 Albert Einstein: Photoelectric Effect • light causes electrons to be emitted from metals • quantized energy transfer causes the emission —E=mc2 He was SOOOOO excited!! 2.2 Atomic Architecture: Electrons & Nuclei Electrons Charge = n (-1.6x10-19 C) Charge e 1.76x1011 C/kg Mass m Conclusions: •Electrons are particles. •Electrons have a mass of 9.1 x 10-19 kg Fig 2-15 Millikan’s Oil Drop Experiment Nucleus Discovered 1911 Ernest Rutherford: small dense positive nucleus nucleus is most of the mass of an atom electrons are in the space around the nucleus Ernest Rutherford (1871-1937) • Won the Nobel Prize in Chemistry in 1908 • “It was quite the most incredible event..... It was almost as if a gunner were to fire a shell at a piece of tissue and the shell bounced right back!!!!! ” 2.2 Atomic Architecture: Electrons & Nuclei The Nucleus • Most of the mass & all of the positive charge of the atom are in the nucleus, which occupies only 1 part in 1014 of the atoms volume. • Electrons occupy a huge volume in comparison to the nucleus, but have relatively small masses. Fig 2-17 Schematic drawing of an atom 2.2 Atomic Architecture: Electrons & Nuclei The Nucleus • Protons account for the positive charge of the nucleus and have a positive charge with a magnitude equal to the negative charge of an electron. Protons have a mass about 2000 times that of an electron. • The mass of a neutrons is about the same as the mass of a proton, but a neutron is electrically neutral. Fig 2-17 Schematic drawing of an atom 2.2 Atomic Architecture: Electrons & Nuclei The Nucleus Table 2-1 Atomic Building Blocks Name Symbol Electron Proton Neutron e p n Charge -1.6022 x10-19 C +1.6022 x10-19 C 0 Mass 9.1091 x10-31 kg 1.6726 x10-27 kg 1.6749 x10-27 kg 2.3 Atomic Diversity An element is identified by the charge of its nucleus Atomic number, Z: nulcear charge, number of protons Atomic Definitions I: Symbols, Isotopes,Numbers A X Z The Nuclear Symbol of the Atom, or Isotope X = Atomic symbol of the element, or element symbol A = The Mass number; A = Z + N Z = The Atomic Number, the Number of Protons in the Nucleus N = The Number of Neutrons in the Nucleus Isotopes = atoms of an element with the same number of protons, but different numbers of Neutrons in the Nucleus 2.3 Atomic Diversity Isotopes Atom with the same number of protons, but different number of neutrons. Mass number A Atomic number Z X Element symbol Example: How many protons, neutrons and electrons do each of the following have? 16 8 O 12 6 C 14 6 C 2.3 Atomic Diversity Isotopes A ZX Examples: 16 O 8 12 C 6 14 6C 8 protons, 8 neutrons, 8 electrons 6 protons, 6 neutrons, 6 electrons 6 protons, 8 neutrons, 6 electrons 2.3 Atomic Diversity Isotopes Courtesy of Sachtleben Chemie GmbH (Paint) Courtesy M. Freeman/PhotoLink/Phoyo Disc (Knee Joint) 46Ti = 8.2% 47Ti = 7.4% 48Ti = 73.8% 49Ti = 5.4% 50Ti = 5.2% 2.3 Atomic Diversity Isotopes Cl Cr Ge Sn Fig 2-18 Natural abundance of the isotopes of Cl, Cr, Ge, Sn Problem: Calculate the abundance of the two Bromine isotopes: 79Br = 78.918336 g/mol and 81Br = 80.91629 g/mol , given that the average mass of Bromine is 79.904 g/mol. Plan: Let the abundance of 79Br = X and of 81Br = Y and X + Y = 1.0 Solution: X(78.918336) + Y(80.91629) = 79.904 X + Y = 1.00 therefore X = 1.00 - Y (1.00 - Y)(78.918336) + Y(80.91629) = 79.904 78.918336 - 78.918336 Y + 80.91629 Y = 79.904 1.997954 Y = 0.985664 or Y = 0.4933 X = 1.00 - Y = 1.00 - 0.4933 = 0.5067 %X = % 79Br = 0.5067 x 100% = 50.67% = 79Br %Y = % 81Br = 0.4933 x 100% = 49.33% = 81Br Pg 55 2.3.2 – 2.3.3 Modern Reassessment of the Atomic Theory 1. All matter is composed of atoms. Although atoms are composed of smaller particles (electrons, protons, and neutrons), the atom is the smallest body that retains the unique identity of the element. 2. Atoms of one element cannot be converted into atoms of another element in a chemical reaction. Elements can only be converted into other elements in Nuclear reactions in which protons are changed. 3. All atoms of an element have the same number of protons and electrons, which determines the chemical behavior of the element. Isotopes of an element differ in the number of neutrons, and thus in mass number, but a sample of the element is treated as though its atoms have an average mass. 4. Compounds are formed by the chemical combination of two or more elements in specific ratios, as originally stated by Dalton. 2.4 Charged Atoms: Ions Formation of Cations Ions: electrically charged atomic or molecular particles Cations: ions with positive charges 2.4 Charged Atoms: Ions Formation of Anions Anions: ions with negative charges Net electrical charge is always conserved 2.4 Charged Atoms: Ions Ionic Compounds A solid containing cations & anions in a balanced whole-number ratio. 2 Na(s) + Cl2(g) 2 NaCl(s) Fig 2-21 Courtesy Michael Watson Fig. 2.19 2.4 Charged Atoms: Ions Salt + water Ionic Solutions Pure water Fig 2-23 Sugar + water Courtesy Ken Karp 2.4 Charged Atoms: Ions Ionic Solutions + - Fig 2-24 2.5 Energy of Atoms and Molecules Fig 2-24 2.5 Energy of Atoms and Molecules Kinetic Energy, Ekinetic: the energy of directed motion of an object. Ekinetic = 1/2 mu2 (2-1) 1 kg m2s-2 = 1 J Thermal Energy: the energy of random motion, translational, rotational and vibrational. The thermal energy of an object is equal to the sum of the kinetic energy of its atoms. 2.5 Energy of Atoms and Molecules Potential Energy: energy that is stored Chemical Energy: potential energy stored as chemical bonds Electrical Energy: potential energy that is the result of electrical forces between charged objects q1q2 Eelectric k r (2-2) Radiant Energy: the energy of electromagnetic radiant (light, photons) 2.5 Energy of Atoms and Molecules Conservation of Energy Energy is neither created nor destroyed in any process only transferred from one body to another, or changed from one form to another.