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Atomic Structure Unit 2 http://www.unit5.org/chemistry Guiding Questions How do we know atoms exist? How do we know that electrons, protons, and neutrons exist? What is radiation and what does it come from? Is radiation safe? Where does matter come from? How are elements formed? Are all atoms of an element the same? How do we measure atoms if they are so small? How do we know what stars are made of? What is wrong with this picture? Atomic Theory and Atomic Structure You should be able to Discuss the development of the atom from its earliest model to the modern day atom. Identify the correct number of subatomic particles for atoms, ions, and isotopes. Calculate the average atomic mass of an atom from isotopic data. Name compounds and write chemical formulas for binary compounds, ternary compounds (those with polyatomic ions), and acids. Memorize the chemical formulas and charged of the polyatomic ions and the most common transition metal ions. Atomic Structure and Periodicity You should be able to Identify characteristics of and perform calculations with frequency and wavelength. Know the relationship between types of electromagnetic radiation and energy; for example, gamma rays are the most damaging. Know what exhibits continuous and line spectra. Know what each of the four quantum numbers n, l, m, and ms represents. Identify the four quantum numbers for an electron in an atom. Write complete and shorthand electron configurations as well as orbital diagrams for an atom or ion of an element. Identify the number and location of the valence electrons in an atom. Apply the trends in atomic properties such as atomic radii, ionization energy, electronegativity, electron affinity, and ionic size. The Hellenic Market Fire ~ Water Earth Air A Brief History of Chemistry • In fourth century B.C., ancient Greeks proposed that matter consisted of fundamental particles called atoms. • Over the next two millennia, major advances in chemistry were achieved by alchemists. Their major goal was to convert certain elements into others by a process called transmutation. Copyright © 2007 Pearson Benjamin Cummings. All rights reserved. The Greeks History of the Atom • Not the history of atom, but the idea of the atom • In 400 B.C the Greeks tried to understand matter (chemicals) and broke them down into earth, wind, fire, and air. ~ ~ • Democritus and Leucippus Greek philosophers Greek Model “To understand the very large, we must understand the very small.” Democritus • Greek philosopher • Idea of ‘democracy’ • Idea of ‘atomos’ – Atomos = ‘indivisible’ – ‘Atom’ is derived • No experiments to support idea • Continuous vs. discontinuous theory of matter Democritus’s model of atom No protons, electrons, or neutrons Solid and INDESTRUCTABLE Democritus DEMOCRITUS (400 BC) – First Atomic Hypothesis Atomos: Greek for “uncuttable”. Chop up a piece of matter until you reach the atomos. Properties of atoms: • indestructible. • changeable, however, into different forms. • an infinite number of kinds so there are an infinite number of elements. • hard substances have rough, prickly atoms that stick together. • liquids have round, smooth atoms that slide over one another. • smell is caused by atoms interacting with the nose – rough atoms hurt. • sleep is caused by atoms escaping the brain. • death – too many escaped or didn’t return. • the heart is the center of anger. • the brain is the center of thought. • the liver is the seat of desire. “Nothing exists but atoms and space, all else is opinion”. Four Element Theory FIRE • Plato was an atomist • Thought all matter was composed of 4 elements: – – – – – Earth (cool, heavy) Water (wet) Fire (hot) Air (light) Ether (close to heaven) Hot Dry ‘MATTER’ AIR Wet EARTH Cold WATER Relation of the four elements and the four qualities Blend these “elements” in different proportions to get all substances Some Early Ideas on Matter Anaxagoras (Greek, born 500 B.C.) –Suggested every substance had its own kind of “seeds” that clustered together to make the substance, much as our atoms cluster to make molecules. Empedocles (Greek, born in Sicily, 490 B.C.) –Suggested there were only four basic seeds – earth, air, fire, and water. The elementary substances (atoms to us) combined in various ways to make everything. Democritus (Thracian, born 470 B.C.) –Actually proposed the word atom (indivisible) because he believed that all matter consisted of such tiny units with voids between, an idea quite similar to our own beliefs. It was rejected by Aristotle and thus lost for 2000 years. Aristotle (Greek, born 384 B.C.) –Added the idea of “qualities” – heat, cold, dryness, moisture – as basic elements which combined as shown in the diagram (previous page). Hot + dry made fire; hot + wet made air, and so on. O’Connor Davis, MacNab, McClellan, CHEMISTRY Experiments and Principles 1982, page 26, Alchemy • After that chemistry was ruled by alchemy. • They believed that that could take any cheap metals and turn them into gold. • Alchemists were almost like magicians. – elixirs, physical immortality Alchemy Alchemical symbols for substances… .. . ...... . ..... GOLD SILVER COPPER IRON SAND transmutation: changing one substance into another D In ordinary chemistry, we cannot transmute elements. Contributions of alchemists: Information about elements - the elements mercury, sulfur, and antimony were discovered - properties of some elements Develop lab apparatus / procedures / experimental techniques - alchemists learned how to prepare acids. - developed several alloys - new glassware Timeline Greeks (Democritus ~450 BC) Discontinuous theory of matter ALCHEMY 400 BC 300 AD Issac Newton (1642 - 1727) 1000 2000 Greeks (Aristotle ~350 BC)) Continuous theory of matter American Independence (1776) Dalton Model of the Atom Late 1700’s - John Dalton- England Teacher- summarized results of his experiments and those of others Combined ideas of elements with that of atoms in Dalton’s Atomic Theory The Atomic Theory of Matter • In 1803, Dalton proposed that elements consist of individual particles called atoms. • His atomic theory of matter contains four hypotheses: 1. All matter is composed of tiny particles called atoms. 2. All atoms of an element are identical in mass and fundamental chemical properties. 3. A chemical compound is a substance that always contains the same atoms in the same ratio. 4. In chemical reactions, atoms from one or more compounds or elements redistribute or rearrange in relation to other atoms to form one or more new compounds. Atoms themselves do not undergo a change of identity in chemical reactions. Copyright © 2007 Pearson Benjamin Cummings. All rights reserved. Foundations of Atomic Theory Law of Conservation of Mass Mass is neither destroyed nor created during ordinary chemical reactions. Law of Definite Proportions 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 source of the compound. Law of Multiple Proportions If two or more different compounds are composed of the same two elements, then the ratio of the masses of the second element combined with a certain mass of the first elements is always a ratio of small whole numbers. Conservation of Atoms 2 H2 + O2 2 H2O John Dalton H H H2 O H O2 + H2 H O H2O O H 2O H H O H H 4 atoms hydrogen 2 atoms oxygen Dorin, Demmin, Gabel, Chemistry The Study of Matter , 3rd Edition, 1990, page 204 4 atoms hydrogen 2 atoms oxygen Conservation of Mass High voltage electrodes Before reaction After reaction glass chamber O2 High voltage H2O O2 5.0 g H2 H2 80 0 g H2 g O2 300 g (mass of chamber) + 385 g total 45 ? g H2O 40 g O2 300 g (mass of chamber) + 385 g total Dorin, Demmin, Gabel, Chemistry The Study of Matter , 3rd Edition, 1990, page 204 Law of Definite Proportions Joseph Louis Proust (1754 – 1826) • Each compound has a specific ratio of elements • It is a ratio by mass • Water is always 8 grams of oxygen for every one gram of hydrogen The Law of Multiple Proportions • Dalton could not use his theory to determine the elemental compositions of chemical compounds because he had no reliable scale of atomic masses. • Dalton’s data led to a general statement known as the law of multiple proportions. • Law states that when two elements form a series of compounds, the ratios of the masses of the second element that are present per gram of the first element can almost always be expressed as the ratios of integers. Copyright © 2007 Pearson Benjamin Cummings. All rights reserved. Daltons Atomic Theory • Dalton stated that elements consisted of tiny particles called atoms • He also called the elements pure substances because all atoms of an element were identical and that in particular they had the same mass. The Atomic Theory of Matter • Dalton’s atomic theory is essentially correct, with four minor modifications: 1. Not all atoms of an element must have precisely the same mass. 2. Atoms of one element can be transformed into another through nuclear reactions. 3. The composition of many solid compounds are somewhat variable. 4. Under certain circumstances, some atoms can be divided (split into smaller particles: i.e. nuclear fission). Copyright © 2007 Pearson Benjamin Cummings. All rights reserved. Dalton’s Atomic Theory 1. All matter consists of tiny particles. Dalton, like the Greeks, called these particles “atoms”. 2. Atoms of one element can neither be subdivided nor changed into atoms of any other element. 3. Atoms can neither be created nor destroyed. 4. All atoms of the same element are identical in mass, size, and other properties. 5. Atoms of one element differ in mass and other properties from atoms of other elements. 6. In compounds, atoms of different elements combine in simple, whole number ratios. Dalton’s Atomic Theory 1. All matter is made of tiny indivisible particles called atoms. 2. Atoms of the same element are identical, those of different atoms are different. 3. Atoms of different elements combine in whole number ratios to form compounds 4. Chemical reactions involve the rearrangement of atoms. No new atoms are created or destroyed. California WEB Structure of Atoms • Scientist began to wonder what an atom was like. • Was it solid throughout with no internal structure or was it made up of smaller, subatomic particles? • It was not until the late 1800’s that evidence became available that atoms were composed of smaller parts. History: On The Human Side 1834 1895 1896 1896 1897 1898 1899 1900 Michael Faraday - electrolysis experiments suggested electrical nature of matter Wilhelm Roentgen - discovered X-rays when cathode rays strike anode Henri Becquerel - discovered "uranic rays" and radioactivity Marie (Marya Sklodowska) and Pierre Curie discovered that radiation is a property of the atom, and not due to chemical reaction. (Marie named this property radioactivity.) Joseph J. Thomson - discovered the electron through Crookes tube experiments Marie and Piere Curie - discovered the radioactive elements polonium and radium Ernest Rutherford - discovered alpha and beta particles Paul Villard - discovered gamma rays 1919 1932 1934 1938 1940 1941 1942 1903 1910 1911 Ernest Rutherford and Frederick Soddy established laws of radioactive decay and transformation Frederick Soddy - proposed the isotope concept to explain the existence of more than one atomic weight of radioelements Ernest Rutherford - used alpha particles to explore gold foil; discovered the nucleus and the proton; proposed the nuclear theory of the atom 1944 1964 Ernest Rutherford - announced the first artificial transmutation of atoms James Chadwick - discovered the neutron by alpha particle bombardment of Beryllium Frederick Joliet and Irene Joliet Curie - produced the first artificial radioisotope Otto Hahn, Fritz Strassmann, Lise Meitner, and Otto Frisch - discovered nuclear fission of uranium-235 by neutron bombardment Edwin M McMillan and Philip Abelson discovered the first transuranium element, neptunium, by neutron irradiation of uranium in a cyclotron Glenn T. Seaborg, Edwin M. McMillan, Joseph W. Kennedy and Arthur C. Wahl - announced discovery of plutonium from beta particle emission of neptunium Enrico Fermi - produced the first nuclear fission chain-reaction Glenn T. Seaborg - proposed a new format for the periodic table to show that a new actinide series of 14 elements would fall below and be analogous to the 14 lanthanide-series elements. Murray Gell-Mann hypothesized that quarks are the fundamental particles that make up all known subatomic particles except leptons. Radioactivity (1896) 1. rays or particles produced by unstable nuclei a. Alpha Rays – helium nucleus b. Beta Part. – high speed electron c. Gamma ray – high energy x-ray 2. Discovered by Becquerel – exposed photographic film 3. Further work by Curies Antoine-Henri Becquerel (1852 - 1908) Radioactivity • One of the pieces of evidence for the fact that atoms are made of smaller particles came from the work of Marie Curie (1876 - 1934). • She discovered radioactivity, the spontaneous disintegration of some elements into smaller pieces. The Effect of an Obstruction on Cathode Rays High voltage source of high voltage shadow cathode yellow-green fluorescence Dorin, Demmin, Gabel, Chemistry The Study of Matter , 3rd Edition, 1990, page 117 Crooke’s Tube - voltage source William Crookes + vacuum tube metal disks magnet A Cathode Ray Tube Source of Electrical Potential Stream of negative particles (electrons) Metal Plate Gas-filled glass tube Zumdahl, Zumdahl, DeCoste, World of Chemistry 2002, page 58 Metal plate Background Information • • • • Cathode Rays Form when high voltage is applied across electrodes in a partially evacuated tube. Originate at the cathode (negative electrode) and move to the anode (positive electrode) Carry energy and can do work Travel in straight lines in the absence of an external field Cathode Ray Experiment 1897 Experimentation • Using a cathode ray tube, Thomson was able to deflect cathode rays with an electrical field. • The rays bent towards the positive pole, indicating that they are negatively charged. The Effect of an Electric Field on Cathode Rays source of high voltage High voltage cathode negative plate _ + anode positive plate Dorin, Demmin, Gabel, Chemistry The Study of Matter , 3rd Edition, 1990, page 117 Cathode Ray Experiment Displacement Volts Anodes / collimators Cathode + Deflection region Drift region Thomson’s Calculations Cathode Ray Experiment • Thomson used magnetic and electric fields to measure and calculate the ratio of the cathode ray’s mass to its charge. Electric = deflection Magnetic = deflection charge of ray particle x charge of ray particle x length of electric x deflection region x field mass of ray velocity of 2 x particle ray particle length of drift region length of magnetic x deflection region x field mass of ray velocity of x particle ray particle length of drift region magnetic deflection electric deflection magnetic field = electric field x velocity Thomson PAPER Conclusions • He compared the value with the mass/ charge ratio for the lightest charged particle. • By comparison, Thomson estimated that the cathode ray particle weighed 1/1000 as much as hydrogen, the lightest atom. • He concluded that atoms do contain subatomic particles - atoms are divisible into smaller particles. • This conclusion contradicted Dalton’s postulate and was not widely accepted by fellow physicists and chemists of his day. • Since any electrode material produces an identical ray, cathode ray particles are present in all types of matter - a universal negatively charged subatomic particle later named the electron J.J. Thomson • He proved that atoms of any element can be made to emit tiny negative particles. • From this he concluded that ALL atoms must contain these negative particles. • He knew that atoms did not have a net negative charge and so there must be balancing the negative charge. J.J. Thomson William Thomson (Lord Kelvin) • In 1910 proposed the Plum Pudding model – Negative electrons were embedded into a positively charged spherical cloud. Zumdahl, Zumdahl, DeCoste, World of Chemistry 2002, page 56 Spherical cloud of Positive charge Electrons Thomson Model of the Atom • J.J. Thomson discovered the electron and knew that electrons could be emitted from matter (1897). • William Thomson proposed that atoms consist of small, negative electrons embedded in a massive, positive sphere. • The electrons were like currants in a plum pudding. • This is called the ‘plum pudding’ model of the atom. - - electrons - - - Rutherford Ernest Rutherford (1871-1937) PAPER • Learned physics in J.J. Thomson’ lab. • Noticed that ‘alpha’ particles were sometime deflected by something in the air. • Gold-foil experiment Animation by Raymond Chang – All rights reserved. Rutherford ‘Scattering’ • In 1909 Rutherford undertook a series of experiments • He fired a (alpha) particles at a very thin sample of gold foil • According to the Thomson model the a particles would only be slightly deflected • Rutherford discovered that they were deflected through large angles and could even be reflected straight back to the source Lead collimator Gold foil a particle source q Rutherford’s Apparatus Rutherford received the 1908 Nobel Prize in Chemistry for his pioneering work in nuclear chemistry. beam of alpha particles radioactive substance circular ZnS - coated fluorescent screen gold foil Dorin, Demmin, Gabel, Chemistry The Study of Matter , 3rd Edition, 1990, page 120 What He Expected • The alpha particles to pass through without changing direction (very much) • Because • The positive charges were spread out evenly. Alone they were not enough to stop the alpha particles California WEB What he expected… Because he thought the mass was evenly distributed in the atom. - - - - What he got… richocheting alpha particles The Predicted Result: expected path expected marks on screen Observed Result: mark on screen likely alpha particle path Interpreting the Observed Deflections . . . . . . beam of alpha particles . . . . . undeflected particles . . . . . gold foil Dorin, Demmin, Gabel, Chemistry The Study of Matter , 3rd Edition, 1990, page 120 . deflected particle Density and the Atom • Since most of the particles went through, the atom was mostly empty. • Because the alpha rays were deflected so much, the positive pieces it was striking were heavy. • Small volume and big mass = big density • This small dense positive area is the nucleus California WEB Rutherford Scattering (cont.) Rutherford interpreted this result by suggesting that the a particles interacted with very small and heavy particles Particle bounces off of atom? Case A Case B Particle goes through atom? Particle attracts to atom? Case C Case D . Particle path is altered as it passes through atom? Table: hypothetical description of alpha particles (based on properties of alpha radiation) observation hypothesis alpha rays don’t diffract ... alpha radiation is a stream of particles alpha rays deflect towards a negatively charged plate and away from a positively charged plate ... alpha particles have a positive charge alpha rays are deflected only slightly by an electric field; a cathode ray passing through the same field is deflected strongly ... alpha particles either have much lower charge or much greater mass than electrons Copyright © 1997-2005 by Fred Senese Explanation of Alpha-Scattering Results Alpha particles Nucleus + + - - + + - + + - + - + - - Plum-pudding atom Nuclear atom Thomson’s model Rutherford’s model Results of foil experiment if plumpudding had been correct. Electrons scattered throughout - + - positive charges + + + + - - + + + - Zumdahl, Zumdahl, DeCoste, World of Chemistry 2002, page 57 - Rutherford’s Gold-Leaf Experiment Conclusions: Atom is mostly empty space Nucleus has (+) charge Electrons float around nucleus Dorin, Demmin, Gabel, Chemistry The Study of Matter , 3rd Edition, 1990, page 120 The Rutherford Atom - - n+ - - - Zumdahl, Zumdahl, DeCoste, World of Chemistry 2002, page 57 This is the modern atom model. Electrons are in constant motion around the nucleus, protons and neutrons jiggle within the nucleus, and quarks jiggle within the protons and neutrons. This picture is quite distorted. If we drew the atom to scale and made protons and neutrons a centimeter in diameter, then the electrons and quarks would be less than the diameter of a hair and the entire atom's diameter would be greater than the length of thirty football fields! 99.999999999999% of an atom's volume is just empty space! Website “The Particle Adventure” Scale of the atom. While an atom is tiny, the nucleus is ten thousand times smaller than the atom and the quarks and electrons are at least ten thousand times smaller than that. We don't know exactly how small quarks and electrons are; they are definitely smaller than 10-18 meters, and they might literally be points, but we do not know. It is also possible that quarks and electrons are not fundamental after all, and will turn out to be made up of other, more fundamental particles. (Oh, will this madness ever end?) Website “The Particle Adventure” Physicists have developed a theory called The Standard Model that explains what the world is and what holds it together. It is a simple and comprehensive theory that explains all the hundreds of particles and complex interactions with only: 6 quarks. 6 leptons. The best-known lepton is the electron. Force carrier particles, like the photon. We will talk about these particles later. All the known matter particles are composites of quarks and leptons, and they interact by exchanging force carrier particles. The Standard Model is a good theory. Experiments have verified its predictions to incredible precision, and all the particles predicted by this theory have been found. But it does not explain everything. For example, gravity is not included in the Standard Model. Website “The Particle Adventure” Discovery of the electron 1807 Davy suggested that electrical forces held compound together. 1833 Faraday related atomic mass and the electricity needed to free an element during electrolysis experiments. 1891 Stoney proposed that electricity exists in units he called electrons. 1897 Thomson first quantitatively measured the properties of electrons. Oil Drop Experiment oil droplets + Small hole Charged plate . . ... ................................... . ........................................ ...... . ........ ........... ...................... ..... ..... . . ........ ..... ............................................. .. oil atomizer ............... . ................................................ . . . . . . . ...... .. . .... . .. . . . .. . . . . .. .......... .. . . . . ...... . .......... .. . ................... ... . .. ...................... . ..................... .. . ...... .......................................... ............................... ... .. . . .. ....... ..... ..... .. ..... ....... ... . . ...... ... . . . . .... .. . .. ..... ... .. .......... .. .. . . .. . .. .. . .... . Charged plate Robert Millikan Telescope oil droplet under observation Bohr Atom The Planetary Model of the Atom Bohr’s Model Nucleus Electron Orbit Energy Levels Bohr Model of Atom Increasing energy of orbits n=3 e- n=2 e- n=1 ee- e- e- ee- e- e- eA photon is emitted with energy E = hf The Bohr model of the atom, like many ideas in the history of science, was at first prompted by and later partially disproved by experimentation. http://en.wikipedia.org/wiki/Category:Chemistry Quantum Mechanical Model Niels Bohr & Albert Einstein Modern atomic theory describes the electronic structure of the atom as the probability of finding electrons within certain regions of space (orbitals). Development of Atomic Models Thomson model Rutherford model In the nineteenth century, Thomson described the atom as a ball of positive charge containing a number of electrons. In the early twentieth century, Rutherford showed that most of an atom's mass is concentrated in a small, positively charged region called the nucleus. Bohr model After Rutherford's discovery, Bohr proposed that electrons travel in definite orbits around the nucleus. Quantum mechanical model Modern atomic theory described the electronic structure of the atom as the probability of finding electrons within certain regions of space. Modern View • The atom is mostly empty space • Two regions – Nucleus • protons and neutrons – Electron cloud • region where you might find an electron Dalton proposes the indivisible unit of an element is the atom. Review Models of the Atom Thomson discovers electrons, believed to reside within a sphere of uniform positive charge (the “plum-pudding model). Rutherford demonstrates the existence of a positively charged nucleus that contains nearly all the mass of an atom. Bohr proposes fixed circular orbits around the nucleus for electrons. In the current model of the atom, electrons occupy regions of space (orbitals) around the nucleus determined by their energies. Copyright © 2007 Pearson Benjamin Cummings. All rights reserved. Models of the Atom "In science, a wrong theory can be valuable and better than no theory at all." - Sir William L. Bragg e e + e + e + + e +e + e e + e + e Dalton’s Greek model model (400 (1803) B.C.) Thomson’s plum-pudding model (1897) Bohr’s model (1913) Dorin, Demmin, Gabel, Chemistry The Study of Matter , 3rd Edition, 1990, page 125 - - + Rutherford’s model (1909) Charge-cloud model (present) Models of the Atom e e + e - + e + + e +e +e e + e + e Dalton’s model Greek model (1803) (400 B.C.) 1803 John Dalton pictures atoms as tiny, indestructible particles, with no internal structure. 1800 - Thomson’s plum-pudding model (1897) - + Rutherford’s model (1909) 1897 J.J. Thomson, a British 1911 New Zealander scientist, discovers the electron, leading to his "plum-pudding" model. He pictures electrons embedded in a sphere of positive electric charge. Ernest Rutherford states that an atom has a dense, positively charged nucleus. Electrons move randomly in the space around the nucleus. 1805 ..................... 1895 1900 1905 1910 1904 Hantaro Nagaoka, a Japanese physicist, suggests that an atom has a central nucleus. Electrons move in orbits like the rings around Saturn. Dorin, Demmin, Gabel, Chemistry The Study of Matter , 3rd Edition, 1990, page 125 1915 Bohr’s model (1913) 1926 Erwin Schrödinger 1913 In Niels Bohr's model, the electrons move in spherical orbits at fixed distances from the nucleus. 1920 1925 Charge-cloud model (present) 1930 develops mathematical equations to describe the motion of electrons in atoms. His work leads to the electron cloud model. 1935 1940 1945 1924 Frenchman Louis 1932 James de Broglie proposes that moving particles like electrons have some properties of waves. Within a few years evidence is collected to support his idea. Chadwick, a British physicist, confirms the existence of neutrons, which have no charge. Atomic nuclei contain neutrons and positively charged protons. Bohr Model Neils Bohr Planetary model After Rutherford’s discovery, Bohr proposed that electrons travel in definite orbits around the nucleus. Bohr’s contributions • Bohr’s contributions to the understanding of atomic structure: 1. Electrons can occupy only certain regions of space, called orbits. 2. Orbits closer to the nucleus are more stable — they are at lower energy levels. 3. Electrons can move from one orbit to another by absorbing or emitting energy, giving rise to characteristic spectra. • Bohr’s model could not explain the spectra of atoms heavier than hydrogen. Copyright © 2007 Pearson Benjamin Cummings. All rights reserved. Particles in the Atom Electrons (-) charge no mass located outside the nucleus 1 amu located inside the nucleus 1 amu located inside the nucleus Protons (+) charge Neutrons no charge Subatomic particles Name Symbol Relative Charge mass Actual mass (g) Electron e- -1 1/1840 9.11 x 10-28 Proton p+ +1 1 1.67 x 10-24 Neutron no 0 1 1.67 x 10-24 Discovery of the Neutron 9 4 Be + 4 2 He 12 6 C + 1 0 n James Chadwick bombarded beryllium-9 with alpha particles, carbon-12 atoms were formed, and neutrons were emitted. Dorin, Demmin, Gabel, Chemistry The Study of Matter 3rd Edition, page 764 *Walter Boethe Humor Two atoms are walking down the street. One atom says to the other, “Hey! I think I lost an electron!” The other says, “Are you sure??” “Yes, I’m positive!” A neutron walks into a restaurant and orders a couple of drinks. As she is about to leave, she asks the waiter how much she owes. The waiter replies, “For you, No Charge!!!” Subatomic Particles ATOM ATOM NUCLEUS NUCLEUS ELECTRONS ELECTRONS PROTONS PROTONS NEUTRONS NEUTRONS POSITIVE Positive CHARGE Charge NEUTRAL Neutral CHARGE Charge NEGATIVE CHARGE Negative Charge equal in a Atomic Most Number of the atom’s mass. neutral atom equals the # of... QUARKS Courtesy Christy Johannesson www.nisd.net/communicationsarts/pages/chem Subatomic Particles • Quarks – component of protons & neutrons – 6 types – 3 quarks = 1 proton or 1 neutron Courtesy Christy Johannesson www.nisd.net/communicationsarts/pages/chem He Bohr - Rutherford diagrams • Putting all this together, we get B-R diagrams • To draw them you must know the # of protons, neutrons, and electrons (2,8,8,2 filling order) • Draw protons (p+), (n0) in circle (i.e. “nucleus”) • Draw electrons around in shells He 2 p+ 2 n0 Li Li shorthand 3 p+ 4 n0 3 p+ 4 n0 2e– 1e– Draw Be, B, Al and shorthand diagrams for O, Na Be B Al 4 p+ 5 n° O 5 p+ 6 n° 13 p+ 14 n° Na 8 p+ 2e– 6e– 8 n° 11 p+ 2e– 8e– 1e– 12 n° Mass Number • mass # = protons + neutrons • always a whole number Neutron + • NOT on the Periodic Table! Electrons Nucleus e- + e- e- + + + + Nucleus e- ee- Carbon-12 Neutrons 6 Protons 6 Electrons 6 Proton Isotopes • Atoms of the same element with different mass numbers. • Nuclear symbol: Mass # 12 Atomic # 6 • Hyphen notation: carbon-12 Courtesy Christy Johannesson www.nisd.net/communicationsarts/pages/chem C Isotopes Neutron + Electrons Nucleus + + + + + Nucleus Proton Proton Nucleus Carbon-12 Neutrons 6 Protons 6 Electrons 6 + + + + Neutron Electrons + + Carbon-14 Neutrons 8 Protons 6 Electrons 6 Nucleus 6Li 7Li 3 p+ 3 n0 3 p+ 4 n0 2e– 1e– 2e– 1e– Neutron Neutron Electrons Electrons + Nucleus + + Nucleus + Nucleus Lithium-6 Neutrons 3 Protons 3 Electrons 3 Proton + + Nucleus Lithium-7 Neutrons 4 Protons 3 Electrons 3 Proton 17 Cl Isotopes 37 • Chlorine-37 – atomic #: 17 – mass #: 37 – # of protons: 17 – # of electrons: 17 – # of neutrons: 20 Courtesy Christy Johannesson www.nisd.net/communicationsarts/pages/chem 37 17 Cl Relative Atomic Mass • 12C atom = 1.992 × 10-23 g • atomic mass unit (amu) • 1 amu = 1/12 the mass of a 12C atom • 1 p = 1.007276 amu 1 n = 1.008665 amu 1 e- = 0.0005486 amu + Electrons Nucleus + Neutron + + + + Nucleus Carbon-12 Neutrons 6 Protons 6 Electrons 6 Proton Average Atomic Mass • weighted average of all isotopes • on the Periodic Table • round to 2 decimal places Avg. (mass)(%) + (mass)(%) Atomic = 100 Mass Courtesy Christy Johannesson www.nisd.net/communicationsarts/pages/chem Average Atomic Mass • EX: Calculate the avg. atomic mass of oxygen if its abundance in nature is 99.76% 16O, 0.04% 17O, and 0.20% 18O. Avg. (16)(99.76) + (17)(0.04) + (18)(0.20) 16.00 Atomic = = amu 100 Mass Courtesy Christy Johannesson www.nisd.net/communicationsarts/pages/chem Average Atomic Mass • EX: Find chlorine’s average atomic mass if approximately 8 of every 10 atoms are chlorine-35 and 2 are chlorine-37. Avg. (35)(8) + (37)(2) Atomic = = 35.40 amu 10 Mass Courtesy Christy Johannesson www.nisd.net/communicationsarts/pages/chem Mass Spectrophotometer magnetic field heaviest ions stream of ions of different masses electron beam gas Dorin, Demmin, Gabel, Chemistry The Study of Matter 3rd Edition, page 138 lightest ions Weighing atoms gas sample enters here . ions accelerate towards charged slit magnetic field deflects lightest ions most filament current ionizes the gas The first mass spectrograph was built in 1919 by F. W. Aston, who received the 1922 Nobel Prize for this accomplishment ions separated by mass expose film • mass spectrometry is used to experimentally determine isotopic masses and abundances • interpreting mass spectra • average atomic weights - computed from isotopic masses and abundances - significant figures of tabulated atomic weights gives some idea of natural variation in isotopic abundances Copyright © 1997-2005 by Fred Senese Mass Spectrometry 198 200 202 Photographic plate 196 - + Stream of positive ions Hill, Petrucci, General Chemistry An Integrated Approach 1999, page 320 199 201 204 Mass spectrum of mercury vapor Mass Spectrum for Mercury (The photographic record has been converted to a scale of relative number of atoms) The percent natural abundances for mercury isotopes are: Hg-196 Hg-198 Hg-199 Hg-200 Hg-201 Hg-202 Hg-204 Relative number of atoms 30 25 198 0.146% 10.02% 16.84% 23.13% 13.22% 29.80% 6.85% 196 200 199 15 10 5 197 198 201 204 Mass spectrum of mercury vapor 20 196 202 199 200 Mass number 201 202 203 204 80 Mercury Isotopes Hg 200.59 The percent natural abundances for mercury isotopes are: A B C D E F G Hg-196 Hg-198 Hg-199 Hg-200 Hg-201 Hg-202 Hg-204 0.146% 10.02% 16.84% 23.13% 13.22% 29.80% 6.85% (% "A")(mass "A") + (% "B")(mass "B") + (% "C")(mass "C") + (% "D")(mass "D") + (% "E")(mass "E") + (% F)(mass F) + (% G)(mass G) = AAM (0.00146)(196) + (0.1002)(198) + (0.1684)(199) + (0.2313)(200) + (0.1322)(201) + (0.2980)(202) + (0.0685)(204) = x 0.28616 + 19.8396 + 33.5116 + 46.2600 + 26.5722 + 60.1960 + 13.974 = x x = 200.63956 amu 92 Separation of Isotopes U 238 Natural uranium, atomic weight = 238.029 g/mol Density is 19 g/cm3. Melting point 1000oC. Two main isotopes: 238 92 235 92 U U 99.3% 0.7% (238 amu) x (0.993) + (235 amu) x (0.007) 236.334 amu + 1.645 amu Because isotopes are chemically identical (same electronic structure), they cannot be separated by chemistry. So Physics separates them by diffusion or centrifuge (mass spectrograph is too slow)… 237.979 amu 17 Average Atomic Mass Cl 35.453 • Assume you have only two atoms of chlorine. • One atom has a mass of 35 amu (Cl-35) • The other atom has a mass of 36 amu (Cl-36) • What is the average mass of these two isotopes? 35.5 amu • Looking at the average atomic mass printed on the periodic table...approximately what percentage is Cl-35 and Cl-36? 55% Cl-35 and 45% Cl-36 is a good approximation Naming Isotopes • Put the mass number after the name of the element • carbon- 12 • carbon -14 • uranium-235 California WEB Calculating averages • You have five rocks, four with a mass of 50 g, and one with a mass of 60 g. What is the average mass of the rocks? • Total mass = (4 x 50) + (1 x 60) = 260 g • Average mass = (4 x 50) + (1 x 60) = 260 g 5 5 • Average mass = 4 x 50 + 1 x 60 = 260 g 5 5 5 California WEB Calculating averages • Average mass = 4 x 50 + 1 x 60 = 260 g 5 5 5 • Average mass = .8 x 50 + .2 x 60 • 80% of the rocks were 50 grams • 20% of the rocks were 60 grams • Average = % as decimal x mass + % as decimal x mass + % as decimal x mass + California WEB Isotopes • Because of the existence of isotopes, the mass of a collection of atoms has an average value. • Average mass = ATOMIC WEIGHT • Boron is 20% B-10 and 80% B-11. That is, B-11 is 80 percent abundant on earth. • For boron atomic weight = 0.20 (10 amu) + 0.80 (11 amu) = 10.8 amu