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Atomic Theory DEMOCRITUS 460 - 370 BC • The Greek philosopher Democritus proposed that all matter was made of small, unbreakable particles he called atoms which means unbreakable. • He believed that atoms were too small to be seen. • Philosophers are not scientists. They do not test their ideas. Instead they use reasoning to back up their beliefs. • To them, human reasoning was superior to experimentation. ARISTOTOLE • The famous philosopher Aristotle, who also lived at that time, argued that all matter was made of only four elements. • For the next two thousand years, Aristotle overshadowed Democritus. • Finally, in the early 1800s, the atomist’s theory was revived by John Dalton. John Dalton 1766-1844 • In 1809, Dalton by proposing the following: a) All matter was made of atoms. b) Atoms were solid spheres. c) Atoms of different elements differed in mass. d) Atoms were indivisible and indestructible. e) Atoms combine to form compounds. J.J. THOMSON 1856-1940 • Before you can understand Thomson’s experiment, 3 properties about electrical charges: a) There are two types of electrical charge: positive and negative. b) Opposite charges attract. c) Like charges repel. • Thomson took Cathod ray tube and added two plates inside the tube and connected them with a wire. • When the plates were not charged, the ray shot straight. Thomson’s Experiment Voltage source - + Vacuum tube Metal Disks Thomson’s Experiment Voltage source Passing + an electric current makes a beam appear to move from the negative to the positive end Thomson’s Experiment Voltage source + By adding an electric field Thomson’s Experiment Voltage source + By adding an electric field he found that the moving pieces were negative Cathod Ray Tube Conclusion • 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 model was called the Plum Pudding Model was named after a popular dessert in England at that time. • It was the first model to propose that smaller charged particles make up the atom. • Thomson’s model lasted less than two decades but it was first to propose the existence of subatomic particles. • In 1911 another scientist who worked in Thomson’s lab improved on his atomic model. ERNEST RUTHERFORD 1871-1937 • One type of radioactivity is when an atom throws out a positively charged particle from the nucleus. • This particle was called an alpha particle (α). • Rutherford used this alpha particle to investigate the structure Rutherford and Geiger in the Cavendish Lab Rutherfold’s Gold Foil Experiement • Uranium is a radioactive element that gives off positive particles (alpha particles). • Rutherford used these positive particles to invest • Rutherford encased uranium in lead (which absorbs alpha particles). • This produced a beam of alpha particles traveling in a straight line. • He fired these positive particles at a thin piece of gold (dense metal). • A screen around the gold to detect where the alpha particles were traveling. Rutherfold’s Gold Foil Experiement • Rutherford shot alpha particles at a thin sheet of gold to observe what happened when the positive α particles passes through the gold atoms. • If Thompson’s model was correct the alpha particles should pass through the diffused positive cloud with ease. Lead block Uranium Fluorescent Screen Gold Foil What he expected Because He thought the mass was evenly distributed in the atom Since he thought the mass was evenly distributed in the atom What he got Rutherfold’s Conclusion • From his observations Rutherford concluded that the atom had a dense, positive central nucleus composed of + charged protons. • He stated that the electrons orbited the nucleus like planets orbiting the Sun. • In 1909 Rutherford proposed his Planetary Model of the Atom. • His model created positively charged protons located in the nucleus and placed electrons in orbit around the nucleus – like planets around the sun. Almost no deflection; few greatly deflected + Checking for understanding Explain Thompson’s conclusions in 3 points: 1. 2. 3. Explain Rutherford’s conclusions in 3 points: 1. 2. 3. Click Below for the Video Lectures Atomic Models History of Atoms Atomic Structure Subatomic Particles • Over the past century scientist have discovered that the atom is composed of 3 subatomic particles: Protons Neutrons Electrons Checking for understanding Draw this diagram. Label all subatomic particles and include their charges. The Proton 1. Symbol = p+ 2. Relative Mass = 1 Atomic Mass Unit (AMU). 3. Actual mass = 1.674 x 10 -24 g 4. Location: Inside the nucleus 5. Electrical charge: Positive. 6. Importance: The atomic number which is the identity of the element. 7. Discovered by: Ernest Rutherford in 1909 Real World Application - PROTON • The electron transport chain, which occurs in the membrane of mitochondria, uses a proton gradient to help produce ATP, a compound our body uses for energy. • Most acidic substances have more free protons (hydrogen ions) in them than hydroxide ions. Vinegar, lemon juice, and hydrochloric acid (HCl) are examples of acidic liquids. • pH is a measure of the number of free protons (hydrogen ions) in a solution. The pH scale ranges from 0-14, with 0 being acidic (more protons) and 14 being basic (fewer protons). pH measurements are widely used to determine the acidity of rain, bodies of water, and liquid waste from factories. • Proton therapy is also a new treatment for treat cancer. A beam of protons is directed towards a tumor and damages the tumor cells' DNA so they cannot reproduce. The Electron 1. Symbol = e2. Relative Mass = 1 /1836 Atomic Mass Unit. 5. Electrical charge: Negative. 3. Actual mass = 6. Importance: The number of 9.11 x 10 -28 g electrons located in the last 4. Location: Energy level energy level determine the outside the nucleus chemical activity of the element. 7. Discovered by: J.J.Thomson in 1897 Real World Application - Electron • • • • • • Microscopes can be made by utilizing properties of electrons. One example is the scanning electron microscope (SEM). By sending a beam of electrons at the surface of an object, a SEM can make images of the surface with up to 500,000 times magnification. SEMs are commonly used to make high resolution images of dead cells, metal surfaces, and fossils. The electron transport chain, which occurs in the membrane of mitochondria, uses proteins to catalyze reduction and oxidation reactions (reactions that exchange electrons between molecules) that produce ATP, a compound our body uses for energy. Electrons moving through a metal wire produce electric current, or electricity. All reduction and oxidation (redox) reactions occur by transferring electrons from one element, ion, or molecule to another. Examples of redox reactions include the formation of salt from elemental sodium and chlorine gas and the corrosion (rusting) of a iron nail in air. Electrochemical cells and batteries produce energy by moving electrons from a cell with an oxidizing reaction to a cell with a reduction reaction. Lasers work by pumping electrons into higher energy level orbitals. When the electrons fall back down into the lower energy orbital, they each release a photon, which we see as light. The Neutron 1. Symbol = n 2. Relative Mass = 1 Atomic Mass Unit (AMU). 3. Actual mass = 5. Electrical charge: Neutral. 1.675 x 10 -24 g 6. Importance: Is responsible for 4. Location: Inside the isotopes (atoms of the same nucleus element with different numbers of neutrons. 7. Discovered by: James Chadwick in 1932 Real World Application - Neutron • • • • Neutron stars can be formed when stars use up all of their fuel. Protons and electrons in the star merge to form neutrons and neutrinos. The neutrons form the neutron star, which is usually around 20 km in diameter, but can be over twice the mass of the sun. Nuclear fission reactions occur when a free neutron hits an atom's nucleus causing it to break apart into two different nuclei, thus forming two different atoms. Some elements, such as uranium-235, not only split into two different atoms when undergoing fission, but also release more neutrons. This allows for a chain reaction to occur as these neutrons go on to hit other uranium atoms and cause them to break apart as well. Nuclear power plants and nuclear weapons work by nuclear fission. Neutrons are used in isotopic labeling, a process where atoms with a larger number of neutrons than usual are placed in a system and tracked to understand where they move in the system. One example of isotopic labeling is labeling atoms in pharmaceuticals to see where they end up in the body. This is a common use for deuterium, which is another stable form of hydrogen. Different isotopes of elements (elements with different numbers of neutrons) are used to date objects. Carbon dating uses the ratio of carbon-14 to carbon-12 to determine the age of organic material up to 60,000 years old. 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) Carbon 6 6 Phosphorus 15 15 Gold 79 79 Mass Number Mass number is the number of protons and neutrons in the nucleus of an isotope. Mass # = 18 Arsenic Phosphorus + p 8 75 16 + 0 n 8 18 33 75 15 31 Isotopes are atoms of the same element Isotopes having different masses due to varying numbers of neutrons. Isotope Protons Electrons Neutrons Hydrogen–1 1 1 0 (protium) Hydrogen-2 (deuterium) 1 1 1 Hydrogen-3 (tritium) 1 1 2 Nucleus Real World Application - ISOTOPES • Isotopes are used in a multitude of everyday objects. Smoke detectors, for instance, often contain a small amount of americium-241. One of the radioactive properties of this material allows for smoke to be detected at an extremely early stage. • Another rising use for radioactive isotopes is food irradiation. This is a process where food is exposed to the radiation of an element, often cobalt-60, though not in direct contact with it. With the high energy particles that are passing through the food, bacteria and microorganisms are killed. Cellular processes that lead to over-ripening and spoiling are also hindered. • Carbon, the main element in organic materials, has a variety of isotopes that are present in living organisms. By analyzing the abundances of these carbon molecules, paleontologists are able to discover the age of organic materials from bones to clothing. Atomic Atomic mass is theMasses average of all the naturally occurring isotopes of that element. Composition % in nature of the nucleus 6 protons 98.89% 6 neutrons Isotope Symbol Carbon12 12C Carbon13 13C 6 protons 7 neutrons 1.11% Carbon14 14C 6 protons 8 neutrons <0.01% Carbon = 12.011 Weight Average Atomic Mass • The atomic masses given on the periodic table are WEIGHT-AVERAGED masses. • This is calculated using both the masses of each isotope and their percent abundances in nature. • For the purposes of simplicity, we will round weight-average mass to the THOUSANDTHS place. • The weight-average mass is based on the abundance of the naturally occurring isotopes of that element Weight Average Atomic Mass • To find the weight-average mass of an element given the mass of each isotope and each isotopes percent abundance: WAM = (massisotope 1 X % ) + (massisotope 2 X % ) + (massisotope 3 X % ) + etc… Atomic Mass Unit (AMU) • amu = atomic mass unit – the ratio of the average mass per atom of the element to 1/12 of the mass of 12C in its nuclear and electronic ground state. • An atomic mass unit is actually an average mass, found by taking the mass of a C-12 nucleus and dividing it by 12 – Hydrogen = 1amu, 1/12 of C Carbon has two stable isotopes Carbon-12 has natural abundance of 98.89% and 12.000 amu Carbon-13 has natural abundance of 1.11% and 13.003 amu Calculate the atomic mass 1. Givens Carbon-12 m=12.000 amu Abundance= 98.89%=0.9889 Carbon-13 m = 13.003 amu Abundance = 1.11%=0.0111 2. Formula atomic mass of carbon-avg = (mass C-12 x nat.abund) + (mass C-13 x nat.abund.) 3. Plug in the #s (12.000amu x 0.9889) + (13.003 amu x 0.0111) = 12.011 amu = 12.0 amu 44 4 Types of Electron Configuration of Elements 1. Shell Configuration • Shows how many electrons are found in each shell (principal energy level). • This is the configuration Niels Bohr would have come up with as the discoverer of the energy level! Shell Number (Principle Electron Level) Number of Electrons to hold 1 2 2 8 3 8 4 18 5 18 6 32 7 32 Shell Configuration (Bohr Diagrams) C 1) Draw a nucleus with the element symbol inside. 2) Carbon is in the 2nd period, so it has two energy levels, or shells. 3) Draw the shells around the nucleus. Shell Configuration (Bohr Diagrams) C 1) Add the electrons. 2) Carbon has 6 electrons. 3) The first shell can only hold 2 electrons. Shell Configuration (Bohr Diagrams) C 1) Since you have 2 electrons already drawn, you need to add 4 more. 2) These go in the 2nd shell. 3) Add one at a time starting on the right side and going counter clock-wise. Shell Configuration (Bohr Diagrams) C 1) Check your work. 2) You should have 6 total electrons for Carbon. 3) Only two electrons can fit in the 1st shell. 4) The 2nd shell can hold up to 8 electrons. 5) The 3rd shell can hold 18, but the elements in the first few periods only use 8 electrons. 2. Sublevel Electron Configuration • Principal energy levels are made up of sublevels, much as a town is made up of streets. • The expanded configuration tells you how many electrons are found in each sublevel of each PEL. • Most of the time (and for all of the configurations you will be responsible for), one sublevel must fill up completely before the next one can get any electrons. Arrangement of Electrons in an Atom Electrons in atoms are arranged as SHELLS (n) SUBSHELLS (l) ORBITALS (ml) row # shell # possibilities are 1-7 7 rows subshell possibilities are s, p, d, or f 4 subshells Each orbital can be assigned no more than 2 electrons! group # # es subshell : 1 orbital , total 2 ep subshell : 3 orbital, total of 6 ed subshell :5 orbital, total of 10 ef subshell: 7 orbital, total of 14 e- s , orbital shapes p orbitals are peanut or dumbbell shaped. d orbitals f orbitals group # = # valence (outside) e- 1A 1 2A 8A 3A 4A 5A 6A 7A 2 Row 3 = 4 # shells 5 6 3B 4B 5B 6B 7B 8B 8B 8B 1B 2B s d p 7 6 7 f Subshells d and f are “special” group # = # valence e- 1A period # = # e- shells 1 8A 2A 3A 4A 5A 6A 7A 2 3 3B 4B 5B 6B 7B 8B 8B 8B 1B 2B 4 3d 5 4d 5d 6d 6 7 d 6 4f 7 5f f Electron Configuration – Spdf notation HELIUM – 2 electrons row # shell # possibilities are 1-7 7 rows 2 Is subshell possibilities are s, p, d, or f 4 subshells group # # valence epossibilities are: s: 1 or 2 p: 1-6 d: 1-10 f: 1-14 Total e- should equal Atomic # 3. Orbital Box Diagram • Shows how many electrons are in each ORBITAL of each sublevel, and what each electron’s SPIN is. • Orbitals are all the same size, they can all fit up to two electrons in them. • The spin of electrons is indicated by arrows up and down. • If the orbital has two electrons in it, the first will have an up spin, and the second will have a down spin. • The number of arrows will equal the number of electrons in the sublevel. Guide to Drawing Orbital Diagrams Drawing Orbital Diagram Draw the orbital diagram for nitrogen. Step 1 Draw boxes to represent the occupied orbitals. Nitrogen has an atomic number of seven, which means it has seven electrons. Draw boxes to represent the 1s, 2s, and 2p orbitals. Drawing Orbital Diagram Step 2 Place a pair of electrons in the last occupied sublevel in separate orbitals. We place the remaining three electrons in the 2s orbitals. Drawing Orbital Diagram Step 3 Place remaining electrons with opposite spins in each filled orbital. First we place a pair of electrons with opposite spins in the 2p orbitals, with arrows in the same direction. Click Below for the Video Lectures Electron Configuration HONORS CHEMISTRY ONLY 3a. Quantum Numbers • Electron energies are addressed in a similar way to a ZIP code. Many addresses in Ulster and northern Orange • county have 125 as the prefix, with the last two digits signifying the actual postal box.\ • For example, New Paltz is 12561, Wallkill is 12589, Newburgh is 12550, Pine Bush is 12566. 3a. Quantum Numbers • There are four identifying characteristics of the energy of a specific electron in an atomic, each more specific than the last. • They are: – n (principal quantum number) = Principal Energy Level (1, 2, 3, 4, etc.) – l (levarotary) = Sublevel (s, p, d, f) – m (magnetic) = Orbital – s (spin) = Spin (+ 1/2, - 1/2) 3a. Quantum Numbers • n , principal quantum number –based on Bohr’s observations of line spectra for different elements –‘n’ relates to the main energy of an electron –allowable values: n = 1, 2, 3, 4, … –electrons with higher ‘n’ values have more energy 3a. Quantum Numbers • l , The Secondary Quantum Number – based on the observation that lines on line spectra are actually groups of multiple, thin lines – ‘l ’ relates to the shape of the electrons’ orbits – allowable values: l = 0 to l = n - 1 • i.e. for n = 4: l = 0, 1, 2, or 3 – the ‘l ’ values 0, 1, 2, and 3 correspond to the shapes we will call s, p, d and f, respectively 3a. Quantum Numbers • ml , the Magnetic Quantum Number – based on the observation that single lines on line spectra split into new lines near a strong magnet – ‘ml ’ relates to the direction/orientation of the electrons’ orbits – allowable values: ml = - l to + l • i.e. for l = 2: ml = -2, -1, 0, 1, or 2 – electrons with the same l value but different ml values have the same energy but different orientations 3a. Quantum Numbers • ms , The Spin Quantum Number – based on the observation that magnets could further split lines in line spectra, and that some elements exhibit paramagnetism – ‘ms ’ relates to the ‘spin’ of an electron – allowable values: ms = - ½ or + ½ • i.e. for any possible set of n, l, and ml values, there are two possible ms values – when two electrons of opposite spin are paired, there is no magnetism observed; an unparied electron is weakly magnetic ms , The Spin Quantum Number HONORS ONLY ENDS HERE Click Below for the Video Lectures Quantum Mechanical Model 4. Electron (Lewis) Dot Diagram • VALENCE ELECTRONS – the electrons in the outermost shell (furthest energy level from the nucleus), which is also called the valence shell. – The number of valence electrons that an atom has can be determined by the last number in the basic electron configuration. The number of valence electrons that an atom has determines its physical and chemical properties Group 1 (alkali metals) have 1 valence electron Group 2 (alkaline earth metals) have 2 valence electrons Group 13 elements have 3 valence electrons Group 14 elements have 4 valence electrons Group 15 elements have 5 valence electrons Group 16 elements have 6 valence electrons Group 17 (halogens) have 7 valence electrons Group 18 (Noble gases) have 8 valence electrons, except helium, which has only 2 Transition metals (“d” block) have 1 or 2 valence electrons Lanthanides and actinides (“f” block) have 1 or 2 valence electrons Lewis Dot Diagram • using dots in groups of 2 around the symbol of the atom to represent the valence electrons. • For every atom, the valence electrons will occupy only s and p orbitals. • The s electrons fill up first, then the p electrons fill, up electrons first, followed by the downs, just like in the box diagram. The Electron Dot diagram for Nitrogen Nitrogen has 5 valence electrons. First we write the symbol. Then add 1 electron at a time to each side. Until they are forced to pair up. N Checking for understanding Draw orbital Draw Lewis diagrams dot diagrams Carbon Helium Fluorine Electrons Atomic theory Overview 1) The Humble Beginnings Democritus (460-370 BC) and Leucippus (~500 BC) • The atom is an indestructible thing, it is the smallest piece that any substance can be broken in to. • It is indivisible, that is, it cannot be broken down any further. 2) Thousands of years passed: John Dalton (1808) • Atoms are the smallest part that any sample of element can be broken into. • Atoms of the same element have the same atomic mass, atoms of different elements have different atomic mass. 3) Not so much time passed: a Crookes Tube in J. J. Thomson (1897) • The atom is a sphere made of a diffuse (thin) positive charge, in which negatively charged electrons are embedded (stuck). • He called his model the “plum pudding” model, but who eats plum pudding anymore? It’s more like a “chocolate chip cookie dough” model, where the atom is a positively charged cookie dough ball with negative chocolate chip electrons stuck in it. 4) But then Ernest Rutherford discovered the alpha particle and HAD to play with it! (1911) • The atom is made of a small, dense, positively charged nucleus with electrons orbiting outside the nucleus at a distance with empty space making up the rest of the atom. • The majority of an atom’s volume is empty space, and the majority of the atom’s mass is in the nucleus. 5) He saw the light! Broken up into bright lines though a spectroscope! Neils Bohr! (1913) • Bohr observed the light given off when several elements are heated and give off light. Different elements gave off different colors of light. • When this light was passed through a prism, the light was broken up into lines of color. Each element’s lines were different. • Bohr figured that electrons falling from high energy levels to low energy levels were causing the light. • Each element’s spectrum of colored lines was different, meaning that the energy levels of different elements have a different amount of energy. • This process, called spectroscopy, is useful for identifying element samples. 6) Werner Heisenberg may have slept here: we’re uncertain! The Quantum-Mechanical Model • The atom contains a small, dense positive nucleus surrounded by electrons that travel in a wave-like motion around the nucleus. • This motion is modified by mass and charge interactions between electrons and the nucleus. • The interactions and the fast speed of the electron make it impossible to know with any certainty both where an electron is and where it is going in any particular instant. • All we can know is the general area of space in which the electron might be found. They very from the most general location to the most specific. • Electrons travel in principal energy levels, which are made up of sublevels, which are made up of orbitals that contain up to two electrons each. If two electrons are in the same orbital, they will spin in opposite directions. 1) Electrons (charged –1 each, with a mass of 1/1836 amu each) surround the nucleus of the atom in distinct energy levels. Electrons occupy the lowest possible energy levels when the atom is in the ground state. 2) When electrons are given energy (in the form of light, heat or electricity), electrons will rise in energy level by the same amount of energy that the electrons were given. The more energy electrons absorb, the higher they rise. This is called the excited state. This is in accordance with the Law of Conservation of Energy, which states that energy cannot be created or destroyed by physical or chemical change. 3) Since electrons are negatively charged, and therefore attracted to the positively charged nucleus, they will eventually fall back to the ground state. As the electrons fall back to the ground state, they release the energy that caused them to rise in the first place. 4) The energy is released in the form of photons. They travel at the fastest theoretical speed possible, 3.00 X 108 m/sec, otherwise known as the speed of light. Photons are, in fact, particles of light. 5) The color of the light is determined by the amount of energy lost by the electron when it dropped back to the ground state. Light particles travel in a wave pattern. The length of each wave is called, strangely enough, a wavelength. The more energy a photon has, the shorter its wavelength is. An excited lithium atom emitting a photon of red light to drop to a lower energy state. An excited H atom returns to a lower energy level. Click Below for the Video Lectures Light and Matter Electromagnetic Spectrum The Nature of Light • The electromagnetic spectrum includes many different types of radiation. • Visible light accounts for only a small part of the spectrum • Other familiar forms include: radio waves, microwaves, X rays • All forms of light travel in waves Copyright McGraw-Hill 2009 109 Electromagnetic Spectrum Figure 06.01Figure 06.01 Copyright McGraw-Hill 2009 110 Wave Characteristics • Wavelength: (lambda) distance between identical points on successive waves…peaks or troughs • Frequency: (nu) number of waves that pass a particular point in one second • Amplitude: the vertical distance from the midline of waves to the top of the peak or the bottom of the trough Copyright McGraw-Hill 2009 111 Copyright McGraw-Hill 2009 112 Wave Characteristics • Wave properties are mathematically related as: c = where c = 2.99792458 x 108 m/s (speed of light) = wavelength (in meters, m) = frequency (reciprocal seconds, s1) Copyright McGraw-Hill 2009 113 Wave Calculation • The wavelength of a laser pointer is reported to be 663 nm. What is the frequency of this light? c = c 9 10 m 7 663 nm 6.63 10 m nm 3.00 108 m/s 14 1 4.52 10 s 7 6.63 10 m Copyright McGraw-Hill 2009 114 • Calculate the wavelength of light, in nm, of light with a frequency of 3.52 x 1014 s-1. c = c 3.00 10 m/s 7 8.52 10 m 14 1 3.52 10 s 8 9 10 nm 7 8.52 10 m 852 nm m Copyright McGraw-Hill 2009 115