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Chemistry Chapter 4 The Structure of the Atoms Top Ten Table 3.1 Table 3.3 History of Chemistry • 400 B.C. Greeks proposed 4 elements – Earth – Fire – Water – Air • Next 2000 years—alchemy • During this period discoveries were made – Hg, S, Sb – prepared acids Controversial Greek Thinking! Democritus (460-370 B.C): -Matter is composed of “atomos” (now atoms) -Atoms were homogeneous & indivisible -Could not answer what holds atoms together Aristotle (384 B.C.-322B.C.): -Matter was continuous and indefinitely divisible (did not believe in atoms) -Matter made of earth, fire, air, & water -Idea was accepted for nearly 2000 years! Indivisible or Divisible? Democritis vs. Aristotle Late 1700’s • Most chemists accepted element definition • Understood elements combined to form compounds with various properties • Disagreed whether compounds are always in the same ratio What happened in 1790? • Study of matter was revolutionized by new emphasis on Quantitative Analysis • Aided by improved balances • Measurements were actually ACCURATE!!! Robert Boyle • Founder of Modern Chemistry (1627-1691) • Took the “Al” out of Alchemy (although he started as one) • First scientist to understand the importance of careful measurement • Insisted science be based on experiments • Famous for P=1/V Antoine Lavoisier • Father of Modern Chemistry (1743-1794) • Recognized and named hydrogen and oxygen • Introduced the metric system • Wrote first list of elements and revised nomenclature • Because of prominence in pre-revolutionary government, was beheaded at the height of the French revolution John Dalton—Beginning of Modern Atomic Theory • Englishman from a Quaker family (1766-1844) • Revolutionized chemistry by emphasizing that atoms can have weights and weights can be measured (quantitative) • Opened a school at age 12 • Color blind/researched • Interested in botany • Theory not accepted until 1905 Albert Einstein paper Dalton’s Atomic Theory (1808) Matter is composed of extremely small particles called atoms Atoms are indivisible and indestructible. Atoms of a given element are identical in size, mass, and chemical properties. Atoms of specific element are different from those of another element. Different atoms combine in simple whole-number ratios to form compounds. In chemical reactions, atoms are separated, combined, or rearranged Dalton vs. Today Matter is composed of extremely small particles called atoms. True Atoms are indivisible and indestructible. Made up of smaller particles (protons, neutrons, & electrons) except in nuclear chemistry. Atoms of a given elemet are Atoms of a given element identical in size, mass, and have same p+ and e-, but may chemical properties. differ in # of neutrons Atoms of a specific element are different from those of another element True, how we identify them Different atoms combine in simple whole-number ratios to form compounds True, Law of Multiple Proportions In chemical reactions, atoms are separated, combined, or True Law of Conservation of Mass Mass is neither created nor destroyed during chemical or physical reactions. Total mass of reactants = Total mass of products Antoine Lavoisier Law of Multiple Proportions • If two or more different compounds are composed of the same two elements, then the ratio is always small whole numbers. (CO, CO2) What does this mean? (Law of Definite Composition) • 50.0 g sample of pure H2O decomposed into its elements – would find 5.6 g H and 44.4 g oxygen – % mass would be: mass H total mass = 5.60 g 50.0 g x 100 = 11.2% H mass 0 total mass = 44.4 g 50.0 g x 100 = 88.8% O Law of Definite (or Constant) Composition • The fact that a chemical compound contains the same elements in exactly the same proportions by mass regardless of the size of the sample or the source of the compound. Figure 3.2: Representation of NO, NO2, and N2O. What does an atom look like? (Sketch it on your paper!) This is The Modern Atomic Model • Atom: The smallest particle of an element that retains the properties of the element • Only seen by STM (Scanning Tunneling microscope) Subatomic Particles Particle Charge Mass (amu) Location Electron (e-) J.J.Thomson 1897 Nobel Prize 1906 Robert Milllikan (1910s) -1 5.486x10-4 amu 9.1 x 10-28 g 1/1840 of H Electron cloud Proton (p+) Thomson/Goldstein1907 Rutherford 1920 +1 1.007 amu 1.673 x 10-24 g Nucleus 0 1.009 1.675 x 10-24 g Nucleus Neutron (no) Chadwick 1932 Nobel Prize 1935 Discovery of the Electron In 1897, J.J. Thomson used a cathode ray tube to deduce the presence of a negatively charged particle. Cathode ray tubes pass electricity through a gas that is contained at a very low pressure. Figure 3.7: Schematic of a cathode ray tube. Some Modern Cathode Ray Tubes Mass of the Electron 1909 – Robert Millikan determines the mass of the electron. The oil drop apparatus Mass of the electron is 9.1 x 10-28 g Conclusions from the Study of the Electron Cathode rays have identical properties regardless of the element used to produce them. All elements must contain identically charged electrons. Atoms are neutral, so there must be positive particles in the atom to balance the negative charge of the electrons Electrons have so little mass that atoms must contain other particles that account for most of the mass Thomson’s Atomic Model Thomson believed that the electrons were like plums embedded in a positively charged “pudding,” thus it was called the “plum pudding” model. Based on the following facts: (1) atoms contain small, negatively charged particles called electrons and (2) the atoms of the element behave as if they have no charge at all Ernest Rutherford • 1871-1937 • Learned physics in JJ Thomson’s lab • Did much work with alpha particles (+ charged part with mass) • Most famous for his GOLD FOIL EXPERIMENT Figure 3.5: Rutherford’s experiment. Try it Yourself! In the following pictures, there is a target hidden by a cloud. To figure out the shape of the target, we shot some beams into the cloud and recorded where the beams came out. Can you figure out the shape of the target? The Answers Target #1 Target #2 Figure 3.3: Plum Pudding model of an atom. Figure 3.6: Results of foil experiment if Plum Pudding model had been correct. Figure 3.6: Actual results. Most of the particles passed right through A few particles were deflected VERY FEW were greatly deflected Rutherford’s Findings “Like howitzer shells bouncing off of tissue paper!” Conclusions: The nucleus is small The nucleus is dense The nucleus is positively charged Disbelievers…. • Albert Einstein when to his grave not totally believing it • According to classical physics, the electron would have collapsed into the nucleus • 1910-1930 began the Quantum Physics Revolution (the physics of atomic and subatomic particles) The Atomic Scale Most of the mass of the atom is in the nucleus (protons and neutrons) Electrons are found outside of the nucleus (the electron cloud) Most of the volume of the atom is empty space “q” is a particle called a “quark” The Quark… Oops…wrong Quark! About Quarks… Protons and neutrons are NOT fundamental particles. Protons are made of two “up” quarks and one “down” quark. Neutrons are made of one “up” quark and two “down” quarks. Quarks are held together by “gluons” Figure 3.9: A nuclear atom viewed in cross section. Atomic Number Atomic number (Z) of an element is the number of protons in the nucleus of each atom of that element. Element # of protons Atomic # (Z) 6 6 Phosphorus 15 15 Gold 79 79 Carbon Isotopes Elements occur in nature as mixtures of isotopes. Isotopes are atoms of the same element that differ in the number of neutrons Figure 3.10: Two isotopes of sodium. Mass Number Mass number is the number of protons and neutrons in the nucleus of an isotope. Mass # = p+ + n0 Nuclide p+ n0 e- Mass # Oxygen - 18 8 10 8 18 Arsenic - 75 33 42 33 75 Phosphorus - 31 15 16 15 31 Atomic Masses Atomic mass is the average of all the naturally isotopes of that element. Carbon = 12.011 Isotope Symbol Composition of the nucleus % in nature Carbon-12 12C 6 protons 6 neutrons 98.89% Carbon-13 13C 6 protons 7 neutrons 1.11% Carbon-14 14C 6 protons 8 neutrons <0.01% Isotopes…Again (must be on the test) Isotopes are atoms of the same element having different masses due to varying numbers of neutrons. Isotope Protons Electrons Neutrons Hydrogen–1 (protium) 1 1 0 Hydrogen-2 (deuterium) 1 1 1 Hydrogen-3 (tritium) 1 1 2 Nucleus Chlorine Practice Problem • Chlorine exists as 2 isotopes in nature. Cl-35 (atomic mass 34.969 amu) has a 75.77% relative abundance. Cl-37 has an atomic mass 36.966 amu. 1. What is the % abundance of the Cl37 isotope? Chlorine Practice Problem • Chlorine exists as 2 isotopes in nature. Cl-35 (atomic mass 34.969 amu) has a 75.77% relative abundance. Cl-37 has an atomic mass 36.966 amu Calculate the atomic mass of Chlorine. Chlorine Practice Problem • Chlorine exists as 2 isotopes in nature. Cl-35 (atomic mass 34.969 amu) has a 75.77% relative abundance. Cl-37 has an atomic mass 36.966 amu. How many times more massive is Cl-37 than Cl-35? Atomic # and Atomic Mass The Periodic Law Chinese Periodic Table Triangular Periodic Table “Mayan” Periodic Table Stowe Periodic Table A Spiral Periodic Table The Year 1860…. • First International Congress of Chemists – 60 to 70 of 113 elements had been discovered – Italian chemist Cannizzaro presented method for measurement of atomic mass that all could agree on. John Newlands (1837-1898) • English chemist that in 1864 proposed an organizational scheme for the elements • Noticed properties repeated periodically every 8 elements • Called “Law of Octaves” – Did not work for all elements – Criticized analogy for being “non-scientific” Dimitri Mendeleev – Writing a book about the same time. – Wanted to include new information of atomic masses – Wanted to find an arrangement for all of the information on the 60-70 elements – Lothar Meyer (German chemist) did the same but Mendeelev published first Modern Russian Table Law of Mendeleev: – Properties of the elements recur in regular cycles (periodically) when the elements are arranged in order of increasing atomic mass. Missing? 14 Si 28.09 ?? 50 Sn 118.71 Named missing element “Ekasilicon” From base word “eka” meanging next in order “Ekasilicon” Predicted Properties Atomic Mass 72 amu Density 5.5 g/cm3 Melting Point 825° C Observed Properties “Ekasilicon” Predicted Properties Observed Properties Atomic Mass 72 amu 72.61 amu Density 5.5 g/cm3 5.32 g/cm3 Melting Point 825° C 938° C “Ekasilicon” Predicted Properties Observed Properties Oxide Formula XO2 GeO2 Chloride Formula XCl4 GeCl4 Mendeleev’s Periodic Table Dmitri Mendeleev Henry Moseley (1887-1915) • After more elements were discovered, order of atomic mass wasn’t working • 1913: Discovered atoms have unique number or protons in nucleus and rearranged table in order of atomic number • Periodic Law: The statement that there is a periodic repetition of chemical and physical properties of elements when they are arranged by increasing atomic number. Group or Family Period The Periodic Table Group or family Period Figure 3.12: Elements classified as metals and nonmetals. Properties of Metals Metals have luster (shiny) when smooth and clean Metals are solid @ room T Metals are good conductors of heat and electricity Metals are malleable Metals are ductile Metals have high tensile strength Located left of stairstep Figure 3.17: Spherical atoms packed closely together. Metalloids: • Mostly brittle solids • Properties between metal and non-metal (semi-conductors) • Found in nature only as compounds • Once obtained as free metals, are stable in the presence of air Properties of Nonmetals Carbon, the graphite in “pencil lead” is a great example of a nonmetallic element. Nonmetals are poor conductors of heat and electricity Nonmetals tend to be brittle and dull looking if solid Many nonmetals are gases at room T Examples of Nonmetals Sulfur, S, was once known as “brimstone” Graphite is not the only pure form of carbon, C. Diamond is also carbon; the color comes from impurities caught within the crystal structure Microspheres of phosphorus, P, a reactive nonmetal Figure 3.14: Nitrogen gas contains N2 molecules. Figure 3.13: A collection of argon atoms. Figure 3.14: Oxygen gas contains O2 molecules. Table 3.5 Valence e- • The e- in an element’s outermost orbital; they determine the chemical properties of the element. • Group 1: 1 valence e• Group 2: 2 valence e• Group 13: 3 valence e• Group 14: 4 valence e• Group 15: 5 valence e• Group 16: 6 valence e• Group 17: 7 valence e• Group 18: 8 valence e- Figure 3.19: The ions formed by selected members of groups 1, 2, 3, 6, and 7. Atomic Radius • Metal: one half the distance between two adjacent atoms in a crystal • Non-metal: half the idstance between nuclei of identical atoms that are chemically bonded together. • Trends: • Atomic radii generally decreases left to right across a period (more e-, same E level) • Atomic radii generally increases down a group (added e- and added E level) Ionic Radius • Ion: An atom or a group of atoms that has a charge (due to loss or gain of e-) – Cations (+ charged ions) • Radii become smaller due to loss of e- and decreased repulsion/often lose E level – Anions (- charged ions) • Radii become larger due to increase in # of e- and repulsion among them (w/o change in proton #) Table of Ion Sizes Octet Rule • Atoms tend to gain, lose, or share e- in order to acquire a full set of 8 valence e-. Electronegativity • The relative ability of an element’s atoms to attract e- in a chemical bond. • Defined in Paulings units – F: most electronegative element – Fr: least electronegative element – Noble Gases: none Trends • E.N. increases across a period • E.N. decreases down a group Ionization E • The E required to remove an e- from a gasseous atom • The lower the I.E, the more easily an atom given away its e-. • Trends – I.E. increases from left to right – I.E. decreases from top to bottom Figure 3.11: The periodic table Figure 3.11: The periodic table Figure 3.20: Pure water does not conduct a current. Figure 3.20: Water containing dissolved salt conducts a current. Radioactivity • Radioactivity: a process in which some substance spontaneously emit radiation • Radiation: The rays and particles emitted by the radioactive materials. • Nuclear Reaction: A change in an atom’s nucleus. • Radioactive Decay: emitted radiation in a spontaneous process • Major Breakthrough: 1890’s. • Only happens in radioactive atoms with unstable nuclei Types of Radiation • Alpha Radiation – Made up of alpha particles – Deflected toward a negatively charged plate • Beta Radiation – Made up of beta particles – Deflected toward a positively charged plate • Gamma Radiation – High E radiation that has no charge and mass – Not deflected by electronic or magnetic fields Allotropes • Different forms of a given element • Different properties b/c different arrangement of atoms • EX: diamond, graphite, Buckminsterfullerene (All Carbon!) • Look in your book—p. 70 Periodic Table with Group Names Periodic Table with Group Names