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Atoms: The Building Blocks of Matter S.MORRIS 2006 http://www.youtube.com/watch?v=Ki_5a15 gJYg&feature=related http://www.youtube.com/watch?v=20Nz16 PJaQo&feature=related ATOMS consist of electrons orbiting around a nucleus. ELECTRONS NUCLEUS contains: a. PROTON B. NEUTRON Nucleons total # of subatomic particles in the nucleus of the atom Atomic Theory Atoms are incredibly small! What we know about them is based on indirect evidence History of Atomic Theory Democritus – 460-371 B.C. – ancient Greek philosopher – believed all matter consisted of extremely small particles that could not be divided – atoms, from Greek word atomos, means “uncut” or “indivisible” Aristotle – believed all matter came from only four elements—earth, air, fire and water John Dalton- an English schoolteacher He proposed the ATOMIC THEORY. John Dalton: THE ATOMIC THEORY (1808) 1. All matter is composed of extremely small particles which cannot be subdivided, created or destroyed. 2. Atoms of a given element are identical in physical and chemical properties. 3. Atoms of different elements have different physical and chemical properties. 4. Atoms of different elements combine in simple whole-number ratios to form chemical compounds. 5. In chemical reactions, atoms are combined, separated, or rearranged, but never created, destroyed, or changed. By the late nineteenth century, most people had accepted Dalton's proposal of 1808 that matter was made of atoms. Solid Sphere Model or Billiard Ball Model proposed by John Dalton Dalton- solid sphere reminds us of a Nilla cracker http://www.learner.org/resources/se ries61.html Problems with Dalton’s Atomic Theory? 1. Matter is composed, indivisible particles 2. All atoms of a particular element are identical 3. Different elements have different atoms 4. Atoms combine in certain whole-number ratios 5. In a chemical reaction, atoms are merely rearranged to form new compounds; they are not created, destroyed, or changed into atoms of any other elements. Scientist: J.J. Thomson What did he discover: ______________ His Experiment: Cathode Ray Tube His findings: Electrons are negatively charged embedded in a positive charge. J.J.Thomson’s Atomic Model Plum Pudding Model Thomson believed that the electrons were like plums embedded in a positively charged “pudding,” thus it was called the “plum pudding” model. J.J. Thomson: Discovery of the Electron His Experiment: Cathode Ray Tube 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. J.J. Thomson’s cathode ray tube Thomson knew that rays must have come from the atoms of the cathode because most of the atoms in the air had been pumped out of the tube. Because the cathode ray came from the negatively charged cathode, Thompson reasoned that the ray was negatively charged. He observed that when a paddle wheel was placed in the path of the rays, the wheel would turn. This observation suggested that the cathode ray consisted on tiny particles that were hitting the paddles of the wheel. His experiments showed the cathode ray consists of particles that have mass and a negative charge. These were called _______. J.J. Thomson’s Model of Atom was like a chocolate chip cookie! Enjoy!! DO NOW: What was Dalton’s model of an atom? Who discovered the electron? What was his model called? Scientist:ERNEST RUTHERFORD What did he discover: ________________ His experiment: GOLD FOIL EXPERIMENT (1900’s) His findings: ERNEST RUTHERFORD Discovery of the Nucleus GOLD FOIL EXPERIMENT (1900’s) Worked with alpha particles (helium nuclei) they have a positive charge. HIS EXPERIMENT: • He fired alpha particles (that are helium nuclei) at a thin sheet of gold foil. • Particle hits on the detecting screen (film) were recorded • rutherfords_gold_foil_expt.ppt The above diagram shows what we would expect the result of Rutherford's experiment to be if the "plum pudding" model of the atom is correct. The above indicates the actual result. Most of the alpha particles are only slightly deflected, as expected, but occasionally one is deflected back towards the source. His observations 1. Most alpha rays passed through with little or no deflection 2. One in 20,000 were strongly deflected (every time it was near the center of the atom) RUTHERFORD ACTIVITY HALLWAY Pennies Rolled Marbles Rolled Marbles Pennies Equal Distance From each other 1. What was the charge near the center of the atom? 2. What structure was near the center of the atom? 3. Why were the alpha particles deflected so strongly? Scientist: Niels Bohr, 1913 Danish physicist What did he discover:Bohr refined Rutherford's idea by adding that the electrons were in orbits around the nucleus. Rather like planets orbiting the sun. With each orbit only able to contain a set number of electrons. Bohr’s Atom 1913 electrons in orbits nucleus Bohr’s Atom HELIUM ATOM _______ _______ _________ + _________ N N + - __________ BOHR MODEL of the ATOM 1. Electrons revolve around the nucleus in specific orbits (shells), or energy levels. PRINCIPAL ENERGY LEVELS Each of the orbits in the Bohr model of the atom has a fixed radius. The greater the radius of an orbit (the farther from the nucleus), the greater the energy of the electrons in the orbit. The orbits or shells of the Bohr model are known as principal energy levels. The main energy level occupied by an electron 1 through 7. The period an element is found in tells us its energy level. 2. An atom has energy levels. Electrons can only exist in these energy levels, not in between. 3. When an atom is in the _____________, the electrons exist in the energy levels closest to the nucleus. GROUND STATE: Click on pic Electrons in the first energy level have the lowest potential energy since they are located closest to the nucleus. 4. If an atom receives, energy, the atom becomes excited and electrons jump to higher energy levels. EXCITED STATE: an atom with higher potential energy than in the ground state because electrons have “jumped” to a higher energy level. • • http://www.crocodile-clips.com/absorb/AC4/sample/LR301_mg.html Click on bohr increase energy • http://www.youtube.com/watch?v=vUzTQ Wn-wfE&feature=related • Atom song • http://www.broadeducation.com/htmlDemo s/AbsorbChem/HistoryAtom/page.htm • http://www.crocodileclips.com/absorb/AC4/sample/LR301_mg. html • Click on bohr increase energy • http://www.classzone.com/books/earth_sci ence/terc/content/investigations/es0501/es 0501page03.cfm • Build an atom QUANTA (or photons) One important feature of the Bohr atom is the idea that electrons can only absorb or release energy in discrete, specific amounts. These amounts, or bundles of energy are called quanta corresponding to differences in energy levels of the shells. Spectral lines • • If high voltage is applied to hydrogen gas confined in a gas tube, called a gas discharge tube, light is emitted. If this light is passed through a prism, a series of bright lines of distinct colors is produced. Bohr reasoned that these different colored bands of light were actually quanta of corresponding energy. These quanta were emitted as electrons of hydrogen atoms returned from their higher levels in the excited state to their lower levels in the ground state. These colors are specific and can be used to identify these metals. (Flame Tests) Spectroscopic analysis of the hydrogen spectrum… …produces a “bright line” spectrum •The device designed to observe the separation of light into its component colors (wavelengths) is called a _________________. •The series of bright lines produced when excited electrons return to their original energy levels is called a bright-line spectrum. •Each element has its own unique set of spectral lines which can therefore be used to identify the elements presence. Wave Mechanical Model/ Orbital Model of the Atom Scientists, influenced by studies of the wave behavior of electrons replaced the Bohr model with one that describes the motion of electrons in terms of probability of their positions within the atom. This model considers the electrons to move freely around the nucleus. These regions of the most probable location of electrons are known as ________________. The ______________________ are the most probable locations of finding a tiny electron in an atom. The Structure of the Atom • ORBITALS = 3D region around the nucleus that indicated the PROBABLE location of an electron. • Electrons with higher potential energy occupy orbitals farther from the nucleus. • THE FURTHER AN ELECTRON IS FROM THE NUCLEUS THE GREATER ITS ENERGY!!! Discoveries about the atom Dalton 1. All matter is composed Of extremely small particles which cannot be subdivided, created or destroyed. 2. Atoms of a given element are identical in physical and chemical properties. 3. Atoms of different elements have different physical and chemical properties. 4. Atoms of different elements combine in simple whole-number ratios to form chemical compounds. 5. In chemical reactions, atoms are combined, separated, or rearranged, but never created, destroyed, or changed. JJ Thomson What did he discover: Electron His Experiment: Cathode Ray Tube His findings: Electrons are negatively charged embedded in a positive charge. Rutherford What did he discover: The Nucleus His experiment: GOLD FOIL EXPERIMENT (1900’s) His findings: The atom is mostly empty space. The nucleus is small. The nucleus is dense. The nucleus is positively charged Niels Bohr •Electrons revolve around the nucleus in specific orbits, or energy levels. • An atom has energy levels. Electrons can only exist in these energy levels, not in between. •When an atom is in the ground state, the electrons exist in the energy levels closest to the nucleus. •GROUND STATE: the lowest energy state of an atom; the electrons occupy energy levels closest to the nucleus. •If an atom receives, energy, the atom becomes excited and electrons jump to higher energy levels. •EXCITED STATE: an atom with higher potential energy than in the ground state because electrons have “jumped” to a higher energy level. Solid Sphere Model Wave Mechanical Model/ Orbital model This model suggested that electrons could be considered waves confined to the space around a nucleus. Electron cloudsregions where electrons are likely to be found Subatomic Particles Particle Electron Proton Neutron Charge Mass (amu) Location Particle symbol Mass of atoms are measured in Atomic Mass Units! 1 amu = 1/12 mass the Carbon-12 The Atomic Scale Most of the mass of the atom is in the _________ (protons and neutrons) _________ are found outside of the nucleus (the electron cloud) Most of the volume of the atom is ___________ Atomic Number •The number of _________ in the nucleus of each atom of that element. •Identifies the element. •Every element has a different atomic number. •Atomic number also equals the number of electrons in the neutral atom 6C Element Oxygen Phosphorus Gold # of protons Atomic # Mass Number •Number of ________ and _________ in the nucleus of an isotope. Mass # = p+ + n0 Element p+ Oxygen 33 Phosphorus n0 e- Mass # If the number of ELECTRONS EQUALS the number of PROTONS, the ATOM is electrically ______________. # electron = # protons In an atom the # electrons = # protons The charge of an atom is neutral. How to find # of neutrons? # Neutrons =Mass # - Atomic # 12 6 C FIND THE NUMBER OF NEUTRONS IN THE FOLLOWING: 1. 2. 3. 4. 5. 6. Sodium Calcium Nitrogen Iron Argon Lithium HYDROGEN • Hydrogen is the simplest of all atoms • All hydrogen atoms have only 1 proton and one electron. • Like many other elements, hydrogen atoms can have different numbers of neutrons. PROTIUM • 99.9885% of all hydrogen atoms have 1 proton and 0 neutrons in their nucleus. DEUTERIUM • 0.0115% of hydrogen atoms have 1 proton and 1 neutron in their nucleus. TRITIUM • Very small amounts of hydrogen atoms have 1 proton and 2 neutrons in their nucleus. • Tritium is radioactive. • All hydrogen atoms have only 1 proton and 1 electron. Isotopes Isotopes- elements with the same number of protons and electrons, BUT different numbers of neutrons. So the ATOMIC NUMBER IS THE ______________ the MASS NUMBER IS ____________________ Elements occur in nature as mixtures of isotopes. How are these atoms different from each other how are they the same? Isotope Hydrogen–1 (protium) Hydrogen-2 (deuterium) Hydrogen-3 (tritium) Protons Electrons Neutrons Nucleus Isotopes & Their Uses Bone scans with radioactive technetium-99. Isotopes & Their Uses The tritium content of ground water is used to discover the source of the water, for example, in municipal water or the source of the steam from a volcano. Atomic Mass Atomic mass is the average mass of all the naturally isotopes of that element. Carbon = 12.011 Isotope Symbol Composition of the nucleus 12C Carbon-12 6 protons 6 neutrons 13C Carbon-13 6 protons 7 neutrons 14C Carbon-14 6 protons 8 neutrons % Abundance in nature 98.89% 1.11% <0.01% ATOMIC STRUCTURE Atomic mass Atomic number 4 2 He number of electrons = number of ________ SUMMARY Atomic Number = number of protons in the nucleus of an atom Mass Number = number of protons and neutrons in the nucleus of an isotope Atomic Mass= weighted average of the masses of the existing isotopes of an element Isotopes = same number of protons, different number of neutrons Same atomic number, different atomic mass Ions Are created when an atom loses of gains one or more electrons, it acquires a charge http://web.visionlearning.com/custom/chemistry/animations/CHE1.3-an-ions.shtml Charge of Ion = More electrons than protons = More protons than electrons = 12 6 C +1 # of protons # of electrons Total charge PRACTICE IONS Ion Li +1 Ni +2 Pb +2 Ca +2 Cs +1 # protons # neutrons # electrons Chemical Symbol Number of protons I Number of electrons Number of neutrons Atom or Ion? 53 35 36 11 55 45 12 atom 78 atom Zr atom 12 14 50 Br 12 69 atom 45 atom Ce atom 1 2 27 25 32 84 80 125 Sc Pb 78 73 68 108 Ni atom 50 71 atom HOW TO CALCULATE ATOMIC MASS: Boron exists as 2 isotopes B-10 or B-11 B % Abundance 19.78% B 80.22% B-10 10 5 B-11 11 5 How to calculate Atomic Mass: STEP 1: Take the Mass # (in amu) of each element and multiply by its Percent Abundance STEP 2: Add all of these values together STEP 3: Divide by 100 Atomic Mass of Boron STEP 1: 10 x 19.78 = 197.8 11 x 80.22 = 882.42 STEP 2: 197.8 + 882.42 STEP 3: 1080.22 100 = 1080.22 = 10.802 amu Calculate the Atomic Mass of Chlorine: % Abundance Chlorine – 35 75.53 Chlorine – 37 24.47 Calculate the Atomic Mass of Silicon: % Abundance Si – 28 92.21 Si – 29 4.70 Si – 30 3.09 Calculate the Atomic Mass of Oxygen: % Abundance O-16 99.762 O-17 0.038 O-18 0.200 ELECTRON CONFIGURATION 1. Shows arrangement of electrons 2. Each atom has a distinct electron configuration. 3. In an atom the number of electrons must equal the atomic number. 4. Principal Energy Levels Maximum # of Electrons in Energy Level 1 2 3 4 5 5.The ground state Electron Configuration is found on the periodic table in the lower left hand corner of each box. 6. Electrons generally occupy energy levels in sequence, beginning with those of lowest energy. 7. No more than eight electrons occupy the outermost principal energy level (except that the 1st can only hold two). Element Al K C Electron Configuration Explanation The configuration listed on the periodic table is the ground state electron configuration. Element He O Na F Si Mg Br Ground State Electron Configuration EXCITED ELECTRON CONFIGURATIONS: • When an atom is excited, electrons jump to a higher sublevel or energy levels. • An example of a transition from a ground state electron configuration to an excited state electron configuration would be: Classify the following as ground state electron configurations or excited state electron configurations. Element ground state electron configurations or excited state electron configuration 1-2 2-8-7-3 2-8-7 2-7-4 2-7-1 2-8-1 2-8-7-2 Element Ground State Electron Configuration Ion Ion’s Electron Configuration Na Na + Mg Mg +2 Fe Fe +3 Al Al +3 Li Li +1 Valence Electrons Electrons that occupy the valence energy level Valence Electrons= found in outer most energy level Na 2-8-1 Cl 2-8-7 Atoms can have a maximum of 8 valence electrons (with the exception of Hydrogen and Helium) Lewis Dot Diagrams (Electron Dot Diagrams) • Represent the arrangement of electrons around the nucleus. • Electrons are the DOTS. • Nucleus is the symbol. • ONLY REPRESENT VALENCE ELECTRONS!! • Fill one side first, then one on each side before you pair electrons. Lewis Dot Diagrams (Electron Dot Diagrams) Na B O Mg Cl Ne Si H N THE END 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 Electromagnetic radiation propagates through space as a wave moving at the speed of light. c = C = speed of light, a constant (3.00 x 108 m/s) = frequency, in units of hertz (hz, sec-1) = wavelength, in meters The energy (E ) of electromagnetic radiation is directly proportional to the frequency () of the radiation. E = h E = Energy, in units of Joules (kg·m2/s2) h = Planck’s constant (6.626 x 10-34 J·s) = frequency, in units of hertz (hz, sec-1) http://www.visionlearning.com/library/module_viewer.php? mid=51 http://web.visionlearning.com/custom/chemistry/animations/CHE1.3-an-ions.shtml Rutherford’s Gold Foil Experiment Rutherford’s Work Atomic Models: J.J. Thompson Passed electricity through an uncharged gas – The gas gave off rays to show it was NEGATIVELY charged – How? – Negative charges must come from inside the atom! Electrons! Some Modern Cathode Ray Tubes • • Chemistry Class Atom Model Project • Build a Bhor atom model of neon using materials of your choice. Your model must be a free-hanging, three-dimensional structure with protons, neutrons, electrons, and electron energy levels clearly visible. Hint: this drawing represents a Bohr diagram of a carbon atom. Assignment: This is an individual project - each student is to build their own atom model outside of class using scrap material available at home. Materials must be non-perishable and appropriate for classroom display. Project Deadline: Scoring: Your atom model will score points based on the following: NO atom models will be accepted after the project deadline. A detailed diagram accompanies the model to act as a key to the parts. The model demonstrates pride in workmanship. The following parts must be shown: • • • • • – – – – correct number of protons correct number of neutrons correct number of electrons correct arrangement of electron energy levels • • • 0 points 40 points 20 points • • • • • 10 points 10 points 10 points 10 points Total of 100 Points • • Dalton "proved" his theory with a number of assumptions, each of which is either factually wrong or was used in a logically inconsistent manner. Nonetheless opposition from critics such as Mach, who never believed in atoms, was largely ignored. Throughout the nineteenth century atomism became an idea that came to dominate thought in a number of fields, including political science, sociology, psychology, biology and more. • • • • • • • Bohr's Model Neils Bohr knew about all of these facts, and in the early part of the century was collaborating with Rutherford. He also knew about the existence of line spectra from chemical elements; a document on this topic may be found here. He was struggling to make sense of all of this. As was common with Bohr when confronted with a puzzle, this struggle was nearly all-consuming. Then in 1913 Bohr, by accident, stumbled across Balmer's numerology for the hydrogen spectrum, and in a flash came up with a workable model of the atom. The model asserts that: The planetary model is correct. When an electron is in an "allowed" orbit it does not radiate. Thus the model simply throws out classical electromagnetic theory. Technical note: an allowed orbit is one in which the electron mass times its speed times the radius of the orbit is equal to a positive integer n times Planck's constant divided by 2 pi. The integer n can be 1, 2, 3, 17, 108, etc. In fact, there are an infinite number of allowed orbits corresponding to the infinite number of positive integers. When an electron absorbs energy from incident electromagnetic radiation, it "quantum jumps" into a higher energy allowed state. This higher energy state corresponds to an allowed orbit with a higher value of the integer n. When an electron is in a higher energy state, it can quantum jump into a lower energy state, one with a smaller value of n, emitting all of its energy as a single photon of electromagnetic energy. Bohr atom