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Atomic Structure Atomic Structure-The BIG Picture Discovery of the components of the atom and subsequent modeling of the atomic structure led to explosive advances in chemistry, medicine, and energy Chemistry • The nature of the chemical bond • New molecule synthesis • Predictions about reactivity • Information about how reactions work • Electronics / computer development • New analytical (measuring) methods • Emergence of the field of Nuclear Chemistry ATOMIC STRUCTURE Energy Medicine • Isotope tracers • New drugs • Cancer treatments • New cell screening methods • Nuclear fission • Nuclear fusion • Power plants • Understanding of the nature of the sun, planets, stars, etc • Weapons A History Lesson • Not only does history help you become an educated person, it helps you understand current theories if you see how they developed… A History Lesson • When you learn this stuff, try to put yourself in the time of the person making the discovery. – What was it like back then? – How did society influence thinking? – Would you have helped to make this discovery? – Do you think another problem of that day was more important? A History Lesson • People were thinking about the ATOM at a time when indoor plumbing and electricity were not available, priorities to work on, or imminently possible! A History Lesson 1. Thousands of years ago, a Greek philosopher Democritus speculated if a piece of matter were divided in half enough times, you’d finally find the smallest piece that could not be subdivided any further A History Lesson a. This little piece of matter would have the same properties as the big piece. b. It was called the ATOM. c. He also believed: – matter could not be created, destroyed, or further divided – matter is mostly empty space A History Lesson • So what made the atom of one type of matter different than another? • The Greek philosophers thought SHAPE (geometry) was the key difference. • The Greeks were great practitioners of Geometry… • I wonder if THAT influenced their thinking on the ATOM? A History Lesson Let’s mark this for later Democritus: Shape A History Lesson- 1800 years later… 2. Dalton’s (1766-1844) experimentation on matter led him to believe: a. All atoms were spherical in shape but differed from one another by mass. • Mass was a big deal in Dalton’s time. • Maybe it influenced his thinking? A History Lesson 2. Dalton’s (1766-1844) experimentation on matter led him to believe: b. All matter is composed of atoms c. All the atoms for a given element were identical*. Atoms of a specific element are different from atoms of another element. d. Atoms cannot be created, divided* or destroyed. e. Atoms could combine to make compounds only in whole number ratios f. In a chemical reaction, atoms are separated, combined or rearranged. *later shown to be incorrect (Make sure you marked them!) A History Lesson • These statements have come to be known as Dalton’s Atomic Theory. • Dalton was right about the MASS part of his theory and how they combine to make compounds *Let’s mark this for later: Dalton: mass A History Lesson • Near the turn of the 20th century, evidence began to emerge that suggested charged subatomic parts made up the atom… A History Lesson 3. Thomson’s (1897) experimentation led to his discovery of a negativelycharged subatomic particle! a. The negatively charged particle is called the electron. b. Discovered while studying electricity with a cathode ray tube Cathode Ray Tube (CRT) A History Lesson • Thomson inferred from his data that the ATOM must be composed of a combination of (+) charged “matrix” with (–) charged particles (electrons) dispersed in it • like raisins or plums in a pudding c. Developed the “plum pudding” model of the atom Progression of the Atomic Model…Discovery of the electron “Plums” or Thomson’s “Plum Pudding” Model • He discovered the electron (link) in 1897 before the nucleus was discovered • Later discoveries invalidated this model “Pudding” J.J. Thomson in Philosophical Magazine, 1904 “... the atoms of the elements consist of a number of negatively electrified corpuscles enclosed in a sphere of uniform positive electrification, ... “ • By the way, the cathode ray tube, which Thomson used to generate and study electrons evolved into the television set (or at least became the central component in TVs). Lets mark this for later: Thomson: electron A History Lesson- Refining the Atom 4. Millikan (1909) used Thomson’s work to determine: a. the electron charge (-) b. the mass of an electron is 9.11x10-28 grams c. his experiment was the oil drop experiment Millikan’s Oil Drop Experiment Millikan later determined the mass and charge of the electron A History Lesson- Refining the Atomic Model 5. Rutherford (1911) proved: a. The (+) in an atom was not spread out but concentrated in a central location called the nucleus. b. The volume of an atom is mostly empty space! Gold Foil Experiment Gold Foil Experiment: The Explanation A History Lesson- Refining the Atomic Model c. Rutherford’s gold foil experiment: i. He bombarded thin, gold foil with heavy, positively (+) charged He (alpha) particles and most passed through the foil. ii. Occasionally, the particle would bounce back (as if it were a tennis ball hitting a brick wall!) iii. Rutherford assumed the deflected particles hit a dense, gold nucleus and could not pass because of the large mass of the gold nucleus and its (+) charge Plum Pudding vs. Gold Foil Experiment A History Lesson- Refining the Atomic Model Rutherford’s famous gold foil experiment: • showed that the positive charge of the atom MUST be concentrated in a tiny, yet heavy volume he called the nucleus • almost ALL of the mass of the atom is in the nucleus • very light electrons surround this nucleus • the volume that an atom occupies is mostly empty space Gold foil animation If a nucleus were as big as you are wide, the edge of its atom (outermost electron orbital) would be over a mile away! About 1.25 miles . • *Let’s mark this for later: Rutherford: proton, nucleus Click A History LessonRefining the Atomic Model 6. Bohr (1913)- suggested electrons must move around in well-defined orbits or energy levels a. His experiments suggested that electrons reside at different energy levels because it took more (or less) energy to knock them loose from an atom *Lets mark this for later: Bohr: planetary orbit of the electrons around the nucleus A History Lesson- Refining the Atomic Model 7. Chadwick (1932) discovered the neutron a. The neutron is a particle with no charge but about the same mass as a proton. Chadwick’s experiment • This history lesson gets us closer to the modern atomic model. • We still need to understand how the electrons behave. • BUT we’ll do that in the next unit… II. The Subatomic Particles A. Protons (p+) 1.Positively-charged subatomic particle 2. Contained in the nucleus 3. Confirmed by Rutherford 4. Mass: 1.67 x 10-24 g (1840x more massive than an electron!) II. The Subatomic Particles B. Neutron (n0) 1.Not charged subatomic particle 2. Contained in the nucleus 3. Discovered by Chadwick 4. Mass: 1.67 x 10-24 g (same as proton) II. The Subatomic Particles C. Electrons (e-) 1. Negatively charged subatomic particle (the charge is equal and opposite to the charge of the proton) 2. Surrounding the nucleus 3. Mass: 9.11 x 10-28 gram 4. Tiny mass but occupies the majority of the volume of the atom II. The Subatomic Particles C. Electrons (e-) 5. Each electron has an “electronic address” • Each resides in a well defined energy level some distance from the nucleus. The further from the nucleus, the higher the energy level. 6. Responsible for chemical bonding III. Current Model of the Atom • Spherically-shaped • Small, dense positively-charged nucleus surrounded by a cloud of negativelycharged electrons • Most of the atom is empty space • >99% of mass is in the nucleus • Very small (there are 6.5 x1021 atoms in a drop of water) • Nucleus is held together by strong nuclear forces • THE STRONG NUCLEAR FORCE is the name given to the attractive force that holds protons and neutrons together in the nucleus. • If you think it is weird that likecharged protons can get together in an atom’s nucleus without flying apart, I don’t blame you! • Don’t worry about the theory of how this force works. • Just remember, the STRONG NUCLEAR FORCE is able to overcome like-charge repulsion and hence atomic nuclei are quite stable. Here’s the atom so far: Nucleus PROTONS NEUTRONS RESPONSIBLE FOR MASS, IDENTITY OF THE ELEMENT valence shell RESPONSIBLE FOR CHEMICAL BONDING core electrons RESPONSIBLE FOR NUCLEAR SHIELDING IV. The Subatomic Particles A. Atomic Number 1. Protons are responsible for the nuclear charge 2. # protons = atomic number + + NUCLEUS + + + + PROTONS: RESPONSIBLE FOR NUCLEAR CHARGE & = ATOMIC NUMBER IV. The Subatomic Particles 3. This is the big (or top) number shown on the periodic table 4. # of protons identifies the element If for some reason the number of protons changes (like a nuclear reaction), the ELEMENT CHANGES! THE NUCLEUS + + + NUCLEUS + + + PROTONS: RESPONSIBLE FOR NUCLEAR CHARGE & = ATOMIC NUMBER NEUTRONS “STRONG NUCLEAR FORCE” HOLDS THE NUCLEUS TOGETHER. IT OVERCOMES THE REPULSIVE FORCE OF “LIKE CHARGES” (REMEMBER COULOMB’S LAW) IV. The Subatomic Particles Practice 1. Determine the number of protons in: a. Fluorine 9 b. Magnesium 12 2. Identify the element: Zinc a. 30 protons Chlorine b. 17 protons Sodium c. 11 p+ Hydrogen d. 1 p+ IV. The Subatomic Particles • Remember, an atom is the smallest particle of an element that retains the identity of that element. • When the number of protons (+) and the number of electrons (-) are the same, the atom is neutral (has no net charge). • An ion is a charged atom or group of atoms bonded together. • An ion can be positive or negative. IONS ARE CHARGED ATOMS. IF AN ATOM GAINS ELECTRONS SO THAT IT HAS MORE ELECTRONS THAN PROTONS, IT IS A NEGATIVELY CHARGED ATOM CALLED AN ANION a negative ion IF AN ATOM LOSES ELECTRONS SO THAT IT HAS FEWER ELECTRONS THAN PROTONS, IT IS A POSITIVELY CHARGED ATOM CALLED AN electron + Ca ion Atom about to become a cation Atom about to become an anion electron Neutral Beryllium 2+ Be Beryllium Ion + - + - + - + - Neutral Nitrogen Nitrogen Ion “Nitride Ion” + - + - + - + - + - + - + - 3N For a neutral atom: p+ = eCharge of an atom or ion: p+ - e- = Charge of Ion Atoms are always neutral Ions have a charge Example What are the charges of: Lithium has 3 p+ and 3 e- 3–3=0 Lithium has 3 p+ and 2 e- 3 – 2 = +1 Oxygen has 8 p+ and 8 e- 8–8=0 Oxygen has 8 p+ and 10 e- 8 – 10 = -2 Charges = Oxidation Numbers IV. The Subatomic Particles B. Mass Number 1. Both neutrons and protons are responsible for nearly all the atom’s mass 2. # protons + # neutrons = mass number Ex: An oxygen atom has 8 protons and 8 neutrons and has a mass number = 16 __ Practice with Atomic number and Mass number # of protons + # neutrons = mass number A carbon atom with 6 protons and 6 neutrons 12 has a mass number = ____ # of protons = atomic number The atomic number of carbon 6 is ___. Number of electrons will equal the number of protons for an atom with NO NET CHARGE Ex 1: How many p+, e- and n0 are in an atom of Neon with a mass # 22? 10 p+ • Neon’s atomic number is 10 ___ • Mass number = protons + neutrons 22 = 10 + neutrons 12 = neutrons • If this atom is electrically10 neutral, protons = electrons ___ e- Ex 2: Determine which element has a mass # of 23 and contains 12 n0. Mass number = protons + neutrons 23 = protons + 12 11 = protons p+ = atomic number the element with 11 protons is Na (sodium) _______________ Practice atomic # mass # # of p+ # of no # of e- Atomic mass symbol 7 14 7 7 7 14.007 N 9 19 9 10 9 18.998 F 19 39 19 20 19 39.098 K 27 59 27 32 27 58.933 Co Practice 2. If 2 protons were removed from the nucleus of an oxygen atom, what nucleus remains? O: 8 p+ - 2 p+ Carbon 6 p+ ________ IV. The Subatomic Particles C. Isotopes 1. All atoms of an element must have the same number of protons 2. BUT they may have a different number of neutrons 3. All atoms of an element must have the same atomic # but can have a different mass #. 4. These are called isotopes. isotopes For Isotopes Parts of the Atom Isotopes of Lithium Lithium – 6 Mass # Atomic # 6 3 Li Lithium - 7 7 3 Li ISOTOPES Atoms in the same element with different MASS NUMBER but identical ATOMIC NUMBER. NUCLEI OF ATOMS IN THE SAME ELEMENT + + + + + UNRAVEL + + + ISOTOPE SYMBOL 12 + 6 + 6 PROTONS 6 NEUTRONS C ISOTOPES Atoms in the same element with different MASS NUMBER but identical ATOMIC NUMBER. NUCLEI OF ATOMS IN THE SAME ELEMENT + + UNRAVEL + + + + + ISOTOPE SYMBOL 12 + + 6 + C 6 PROTONS 6 NEUTRONS + + + + + + 13 + + 6 + + 6 PROTONS 7 NEUTRONS C A Z X X –ELEMENT SYMBOL A-MASS NUMBER Z-ATOMIC NUMBER Weighted Averages • The atomic mass of one atom is TINY if reported in kilograms or even grams. • Atomic mass is reported in atomic mass units (amu). • 1 amu is about the mass of one proton or neutron (about 1.67 x 10-24) • Generally, atoms of a given element will have 1 isotope in high abundance and several others in much smaller numbers. • To calculate the average atomic mass, we use a WEIGHTED AVERAGE that accounts for the abundance of each isotope. Ex 1: • Carbon-12 makes up 98.89% of naturallyoccurring carbon. Carbon-13 makes up 1.11% of naturally occurring carbon. Use this information to determine the average atomic mass of carbon. Average atomic mass = (12 X 0.9889) + (13 X 0.0111) = 12.0111 amu Ex 2: • Chromium has 4 naturally-occurring isotopes. Their abundance is as follows: Cr-50 – 4.35%, Cr-52 – 83.79%, Cr-53 – 9.50%, and Cr-54 – 2.36%. Determine the average atomic mass for chromium. (50 X .0435) + (52 X .8379) + (53 X .0950) + (54 X .0236) = 52.0552 amu Practice: 1. The element copper is found to contain 69.1% of copper-63 and 30.9% of copper 65. Calculate the average atomic mass of copper. 63.618 amu 2. Gallium occurs in nature as a mixture of two isotopes. They are gallium-69 with an 60.108% abundance and gallium-71 with a 39.892% abundance. Calculate the atomic mass of gallium. 69.7978 amu 3. Neon has 2 isotopes: neon-20 and neon22. Use the information from the periodic table to determine which occurs in greater abundance. Neon – 20 Because the atomic mass on the periodic table rounds to 20 and not 22 4. Use the table below to calculate the atomic mass of Element X and then identify it. Isotope % Abundance 16X 99.762 17X .038 18X .2 (16 X .99762) + (17 X .00038) + (18 X .002) = 16.0044 amu Oxygen