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Atoms and Elements Mr. Hollister Holliday Legacy High School Regular & Honors Chemistry Experiencing Atoms • Atoms are incredibly small, yet they compose everything. • Atoms are the pieces of elements. • Properties of the atoms determine the properties of the elements. 2 Early Theories of Matter • The ancient Greeks believed that all matter was composed of only four elements: earth, air, water, and fire. 3 First Atomic Theory • Democritus (c.a. 460 BC) was the first thinker ever to reason that matter was not infinitely divisible. • He called the smallest piece of matter ‘atomon’, which in Greek means ‘not cutable’. Democritus 4 The Ancient Greeks • Aristotle (c.a. 380 BC) reasoned, however, that matter was infinitely divisible. • Aristotle’s ideas agreed more closely with those of the Christian church, therefore, they were accepted as fact for over two millenia. Aristotle Enlightenment Thinking • Robert Boyle (c.a. Robert Boyle 1661) wrote a book entitled The Sceptical Chymist. In it, Boyle argues that the four classical elements of the Greeks were actually not elements at all. • He also argued that matter was composed of indivisible particles or atoms. 6 The First Modern Atomic Theory • After two hundred years of experimental results, the first modern atomic theory was developed by John Dalton (c.a. 1805). • Dalton reasoned that indivisible, spherical particles made up all of matter. He also argued that these particles were rearranged during a chemical reaction. John Dalton 7 Dalton’s Original Theory of the Atom and of Reactions 8 Dalton’s Atomic Theory 1. Each Element is composed of tiny, indestructible particles called atoms. Tiny, hard, indivisible, spheres. 2. All atoms of an element are identical. They have the same mass, volume, and other physical and chemical properties. So, atoms of different elements are different. Every carbon atom is identical to every other carbon atom. They have the same chemical and physical properties. However, carbon atoms are different from sulfur atoms. They have different chemical and physical properties. 9 Dalton’s Atomic Theory 3. Atoms combine in simple, whole-number ratios to form molecules of compounds. Because atoms are unbreakable, they must combine as whole atoms. The nature of the atom determines the ratios in which it combines. Each molecule of a compound contains the exact same types and numbers of atoms. Chemical formulas 10 Discovery of the Subatomic Particles • J. J. Thompson (c.a. 1897) performed experiments involving cathode rays that proved that atoms could be broken down into smaller particles. • Thompson called the particle that he discovered an ‘electron’, which means unit of electricity in Greek. J. J. Thomson 11 Cathode Ray Tube (Discovery of the Electron) 12 Thomson’s Results • the cathode rays are made of tiny particles • these particles have a negative charge because the beam always deflected toward the + plate • the amount of deflection was related to two factors, the charge and mass of the particles • every material tested contained these same particles • Thompson called these particles ‘electrons’, because they were the components of all electricity 13 Discovery of the Nucleus • Ernest Rutherford (c.a. 1908) utilized a radioactive (α particles) piece of gold foil to prove that the plumpudding model of the atom was wrong. • In so doing, he discovered the atomic nucleus and the proton. Ernest Rutherford 14 Rutherford’s Conclusions • Atoms were mostly empty space because almost all the particles went straight through • Atoms contain a dense particle that were small in volume compared to the atom but large in mass because of the few particles that bounced back • This dense particle were positively charged because of the large deflections of some of the particles 15 Rutherford’s Interpretation – the Nuclear Model 1) The atom contains a tiny dense center called the nucleus the amount of space taken by the nucleus is only about 1/10 trillionth the volume of the atom 2) The nucleus has essentially the entire mass of the atom the electrons weigh so little they give practically no mass to the atom 3) The nucleus is positively charged the amount of positive charge balances the negative charge of the electrons 4) The electrons are dispersed in the empty space of the atom surrounding the nucleus (electron cloud) 16 Structure of the Atom • Rutherford proposed that the nucleus had a particle that had the same amount of charge as an electron but opposite sign based on measurements of the nuclear charge of the elements • these particles are called protons charge = +1.60 x 1019 C mass = 1.67262 x 10-24 g • since protons and electrons have the same amount of charge, for the atom to be neutral there must be equal numbers of protons and electrons 17 The Nuclear Atomic Model 18 Rutherford’s Model of the Atom atomic radius ~ 100 pm = 1 x 10-10 m nuclear radius ~ 5 x 10-3 pm = 5 x 10-15 m “If the atom is the Cowboys Stadium, then the nucleus is a marble on the 50-yard line.” 19 2.2 Atomic Theory from 1800-1911 20 Some Problems • How could beryllium have 4 protons stuck together in the nucleus? shouldn’t they repel each other? • If a beryllium atom has 4 protons, then it should weigh 4 amu; but it actually weighs 9.01 amu! Where is the extra mass coming from? each proton weighs 1 amu remember, the electron’s mass is only about 0.00055 amu and Be has only 4 electrons – it can’t account for the extra 5 amu of mass 21 There Must Be Something Else There! • to answer these questions, Rutherford proposed that there was another particle in the nucleus – it is called a neutron • neutrons have no charge and a mass of 1 amu mass = 1.67493 x 10-24 g slightly heavier than a proton no charge 22 The Sub-Atomic Particles Subatomic Mass Mass Location Charge Symbol Particle g amu in atom Proton 1.67262 1.00727 nucleus +1 p, p+, H+ empty -1 e, e- 0 n, n0 x 10-24 Electron 0.00091 0.00055 x 10-24 Neutron 1.67493 1.00866 space nucleus x 10-24 23 Elements • each element has a unique number of protons in its nucleus the number of protons define the element • the number of protons in the nucleus of an atom is called the atomic number the elements are arranged on the Periodic Table in order of their atomic numbers • Each element also has a corresponding mass number. Protons + neutrons. 24 25 Elements • Each element has a unique name and symbol. The symbol is either one or two letters One capital letter or one capital letter + one lower case letter. H = Hydrogen = “water-former” Br = Bromine = ‘stench’ Liquid Bromine 26 The Periodic Table of Elements Atomic number Element symbol Atomic mass 27 Elemental & Isotope Notation Atomic number (Z) = number of protons in nucleus Mass number (A) = number of protons + number of neutrons = atomic number (Z) + number of neutrons Mass Number Atomic Number A ZX 238 92 Element Symbol U Isotope Notation = Uranium-238 2.3 Example: How many protons, electrons, and neutrons are in an atom of 52 24 Cr ? Given: Find: Concept Plan: 52 24 Cr therefore A = 52, Z = 24 # p+, # e-, # n0 symbol symbol Relationships: Solution: Check: atomic number atomic & mass numbers # p+ # e- # n0 in neutral atom, # p+ = # emass number = # p+ + # n0 Z = 24 = # p+ # e- = # p+ = 24 A = Z + # n0 52 = 24 + # n0 28 = # n0 for most stable isotopes, n0 > p+ 29 Isotopes: • Isotopes = atoms of the same element with different masses. • Isotopes have different numbers of neutrons. 11 C 6 12 C 6 13 C 6 14 C 6 Carbon-11 Carbon-12 Carbon-13 Carbon-14 30 Neon Number of protons Symbol Number of neutrons A, mass number Percent natural abundance Ne-20 or 20 Ne 10 10 10 20 90.48% Ne-21 or 21 Ne 10 10 11 21 0.27% Ne-22 or 22 Ne 10 10 12 22 9.25% 31 Mass Spectrometer 32 Tro: Chemistry: A Molecular Approach, 2/e Mass Spectrum • A mass spectrum is a graph that gives the relative mass and relative abundance of each particle • Relative mass of the particle is plotted in the x-axis • Relative abundance of the particle is plotted in the y-axis 33 Atomic Mass Atomic mass is the mass of an atom in atomic mass units (amu) By definition: 1 atom 12C “weighs” 12 amu 1 amu = 1.6606 x 10-24 g On this scale 1H = 1.008 amu 16O = 16.00 amu 3.1 Average Atomic Mass • Because in the real world we use large amounts of atoms and molecules, we use average masses in calculations. • Average atomic mass is calculated from the isotopes of an element weighted by their relative abundances. 35 Natural lithium is: 6Li (6.015 amu) = 7.42% 7Li (7.016 amu) = 92.58% Average atomic mass of lithium: (7.42 x 6.015) + (92.58 x 7.016) / 100 = 6.941 amu 36 Charged Atoms • The number of protons determines the element. All sodium atoms have 11 protons in the nucleus. • In a chemical change (aka: a reaction), the number of protons in the nucleus of the atom doesn’t change, however the number of electrons may. • Atoms in a compound that gain or lose electrons and have a charge, these are called ions. 37 Types of Ions Cation – ion with a positive charge If a neutral atom loses one or more electrons it becomes a cation. Na 11 protons 11 electrons Na+ 11 protons 10 electrons Cl- 17 protons 18 electrons Anion – ion with a negative charge If a neutral atom gains one or more electrons it becomes an anion. Cl 17 protons 17 electrons 2.5 Ions of Hydrogen 39 Ions A monatomic ion contains only one atom Na+, Cl-, Ca2+, O2-, Al3+, N3- A polyatomic ion contains more than one atom OH-, CN-, NH4+, NO3- Nitrate Polyatomic Ion 40 Practice—Fill in the Table. Ion p+ e- - Cl K+ 2- S Sr 2+ 41 Practice—Fill in the Table, Continued. + - p e 1- 17 18 1+ 19 18 S2- 16 18 38 36 Ion Cl K Sr 2+ 42 Example —Find the Number of Protons and Electrons in Ca2+. Review • What is the atomic number of boron, B? • What is the atomic mass of silicon, Si? • How many protons does a chlorine atom have? • How many electrons does a neutral neon atom have? • Will an atom with 6 protons, 6 neutrons and 6 electrons be electrically neutral? • Will an atom with 27 protons, 32 neutrons, and 27 electrons be electrically neutral? • Will an Na atom with 10 electrons be electrically neutral? 44