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Chapter 3 Atoms Chapter 3 Section 1 The Atom: From Philosophical Idea to Scientific Theory Objectives • Explain the law of conservation of mass, the law of definite proportions, and the law of multiple proportions. • Summarize the five essential points of Dalton’s atomic theory. • Explain the relationship between Dalton’s atomic theory and the law of conservation of mass, the law of definite proportions, and the law of multiple proportions. Section 3.1 Defining the Atom Around 460 B.C. Greek philosopher, Democritus, developed the idea of atoms. He asked this question: If you break a piece of matter in half, and then break it in half again, how many breaks will you have to make before you can break it no further? Democritus thought that it ended at some point, a smallest possible bit of matter. He called these basic matter particles, atoms. Democritus’s Atomic Philosophy Atoms were indivisible and indestructible Atom comes from the Greek word atomos, meaning “indivisible” His ideas: were based on philosophy didn’t explain chemical behavior lacked experimental support Three Atomic Laws • Rules that explain how matter interacts in chemical reactions: 1790’s- improved balances revolutionized quantitative analysis of chemical reactions.. allowed accurate measure of masses of elements and compounds 1. Law of conservation of mass mass is neither created nor destroyed during ordinary chemical reactions or physical changes Three Atomic Laws • Rules that explain how matter interacts in chemical reactions: 2. Law of definite proportions a chemical compound contains the same elements in exactly the same proportions by mass regardless of the size of the sample or source 3. Law of multiple proportions 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 element is always a ratio of small whole numbers Chapter 3 Section 1 The Atom: From Philosophical Idea to Scientific Theory Law of Conservation of Mass Chapter 3 Section 1 The Atom: From Philosophical Idea to Scientific Theory Law of Multiple Proportions John Dalton (1766-1844) Dalton transformed Democritus’s ideas on atoms into a scientific theory. Dalton used experimental methods not just philosophy Studied ratio in which elements combine in chemical reactions Developed an atomic theory based on results of experiments Note: Dalton was always very smart and at the age of twelve he took over a Quaker school in Cumberland, England. Designated much of his life to researching color blindness because both he and his brother suffered from it. Dalton’s Elements and Compounds Dalton’s Atomic Theory (p. 64) 1. All elements are composed of extremely small particles called atoms. 2. Atoms of an element are identical in size, mass and other properties; atoms of different elements differ in size, mass, and other properties. Dalton’s Atomic Theory (p. 64) 3. Atoms cannot be subdivided, created, or destroyed (law of conservation of mass) 4. Atoms of different elements combine in simple whole-number ratios to form chemical compounds. (law of definite proportions and multiple proportions) 5. In chemical reactions, atoms are combined, separated, or rearranged. (law of conservation of mass) Scanning Tunneling Microscope (STM) Gerd Binnig Dr. Heinrich Rohrer • Invented in 1981 by Gerd Binnig and Heinrich Rohrer working for IBM in Zurich, Switzerland. • Won 1986 Nobel Prize in physics for their discovery • open door to field of Nanotechnology Scanning Tunneling Microscopy Use of Scanning Tunneling Microscopy by IBM Moving Atoms: Making the World's Smallest Movie Scanning Tunneling Microscope (STM) STM provides ability to look at individual atoms detects the electrons of surface atoms maps the position and spacing of atoms Modern Atomic Theory Not all aspects of Dalton’s atomic theory have proven to be correct. We now know that: Atoms are divisible into even smaller particles. A given element can have atoms with different masses (isotopes). Some important concepts remain unchanged: All matter is composed of atoms. Atoms of any one element differ in properties from atoms of another element. Dalton’s Atomic Theory has not been discarded.. only modified! Chapter 3 Atoms Chapter 3 Section 2 The Structure of the Atom Objectives • Summarize the observed properties of cathode rays that led to the discovery of the electron. • Summarize the experiment carried out by Rutherford and his co-workers that led to the discovery of the nucleus. • List the properties of protons, neutrons, and electrons. • Define atom. • Describe nuclear forces Section 3.2 Structure of the Atom One change to Dalton’s atomic theory - atoms ARE divisible Into subatomic particles: Electrons, protons, and neutrons Atom – smallest particle of an element that retains the chemical properties of that element Discovery of Electron- J.J. Thomson p. 69 1897 - J.J. Thomson used cathode ray tube to show presence of negatively charged particles...electrons Note: Called the negatively charged particles “corpuscles” Thomson’s Atomic Model J. J. Thomson cathode ray experiment showed presence of electrons “plum pudding” model - believed e- like plums in + charged “pudding,” charge-to-mass ratio of an electron Joseph John “J.J.” Thomson 1856-1940 Famous for discovery of the first subatomic particle, the electron, and for work on the atomic model. It was soon superseded by his student Ernest Rutherford’s nuclear model. 1906 received Nobel Prize for physics for his research into the electrical conductivity of gases. Thomson was a great teacher and an outstanding scientist. Seven of his students and assistants also received Nobel Prizes for their work. Cathode Ray Tube (CRT) Cathode ray tube (CRT) is a vacuum tube containing an electron gun (source of electrons) and fluorescent screen used to view images. It accelerates and deflects the electron beam onto fluorescent screen to create images. Can be found in everyday devises such as televisions, video games, computers, video cameras, monitors, automated teller machines, oscilloscopes, and radar displays. CRTs have been superseded by more modern display technologies such as LCD, plasma display, and OLED, which as of 2012 offer lower manufacturing and distribution costs. CRT - Cathode Ray Tubes Television Computer Monitor CRT’s pass electricity thru gas contained vacuum tube Charge and Mass of Electron- Robert Millikan Charge of the electron is 1.592 × 10−19 coulomb The oil drop apparatus determined the actual e- charge: 1.592 × 10−19 coulomb) (Somewhat lower than the modern of 1.602 × 10 −19 coulomb) determined actual e- mass: 1/1837 the mass of H atom (9.11 x 10-28 g) Robert Millikan's Oil drop experiment Robert Millikan (1868-1953) Doctorate in Physics from Columbia University 1923 Nobel Prize for Physics Determined charge on the electron as well as Planck’s constant Millikan set up his experiment by spraying oil drops, which fell between positively charged plate and negatively charged plate. He balanced the electric charge with electron's weight, as gravity pulled down on the oil drop, and the electric charges attracted and repelled it. Using this method, he discovered that the charge of an electron is negative. The weight of each drop was determined by observing it's rate of free fall through the air. Originally, he sprayed droplets of water through an electrical field that was intended to suspend them in midair, but the water droplets evaporated too fast, so Millikan switched to using oil. Robert Millikan Starting in 1908, while a professor at the University of Chicago, Millikan worked on an oil-drop experiment in which he measured the charge on a single electron. J.J. Thomason had already discovered the charge-to-mass ratio of the electron. However, the actual charge and mass values were unknown. Therefore, if one of these two values were to be discovered, the other could easily be calculated. Millikan and his then graduate student Harvey Fletcher used the oil-drop experiment to measure the charge of the electron (as well as the electron mass, and Avogadro’s number, since their relation to the electron charge was known). Millikan showed that the results could be explained as integer multiples of a common value (1.592 × 10−19 coulomb), the charge on a single electron. That this is somewhat lower than the modern value of 1.602 176 53(14) x 10−19 coulomb is probably due to Millikan's use of an inaccurate value for the viscosity of air. Conclusions from Electron study: Because atoms are electrically neutral, they must contain a positive charge to balance the negative e- Because electrons have so much less mass than atoms, atoms must contain other particles that account for most of their mass. Discovery of Protons: Eugen Goldstein 1886 - Eugen Goldstein observed + particles (now called “proton”) particles with +charge & relative mass of 1 (or 1840x e-) called them “Canal rays” (anode rays) Eugen Goldstein (1850-1930) Credited by many for the discovery of the existence of protons discoverer of anode rays (canal rays) Worked in Berlin and Potsdam Observatories Ernest Rutherford’s 1911- Gold Foil Experiment Alpha particles fired at gold foil Particles hitting detecting screen were recorded Discovery of Nucleus: Ernest Rutherford most particles passed thru “It was as if you fired a 15-in few deflected [artillery] shell at a piece of VERY FEW greatly deflected tissue paper and it came back and hit you.” Conclusions: a) small, dense nucleus b) + charged nucleus Rutherford’s Atomic Model Experimental evidence: atom mostly empty space almost all mass in nucleus e- distributed around nucleus…occupy most of the volume called “nuclear model” Rutherford’s Atom 3:08 Discovery of Neutrons: James Chadwick 1932 – James Chadwick confirmed existence of the “neutron” particle with no charge n0 mass = p+ mass 1935 - Nobel Prize for Physics Prepared the way for the fission of Uranium-235 which led to developing atomic bomb (Manhattan Project) Atomic Scientists • Democritus named the atom • John Dalton proposed the Atomic Theory • J.J. Thomson’s cathode-ray tube experiments measured the charge-to-mass ratio of an electron. (plum-pudding model) • Robert A. Millikan’s oil drop experiment measured the charge of an electron With this information, scientists were able to determine the mass of an electron. Atomic Scientists • Rutherford’s gold foil experiment led to the discovery of the nucleus.. a very densely packed bundle of matter with a positive electric charge. • James Chadwick proved existence of neutrons.. no charge Subatomic Particles Particle Electron (e-) Proton (p+) Neutron (no) Charge Mass Number Actual Mass (g) Location Outside Nucleus 1- 0 9.11 x 10-28 1+ 1 1.67 x 10-24 Nucleus 0 1 1.67 x 10-24 Nucleus (Electron cloud) Chapter 3 Section 2 The Structure of the Atom The Structure of the Atom An atom is the smallest particle of an element that retains the chemical properties of that element. The nucleus is a very small region located at the center of an atom. The nucleus is made up of at least one positively charged particle called a proton and usually one or more neutral particles called neutrons. Chapter 3 Section 2 The Structure of the Atom The Structure of the Atom Surrounding the nucleus is a region occupied by negatively charged particles called electrons. • Protons, neutrons, and electrons are often referred to as subatomic particles. Chapter 3 Section 2 The Structure of the Atom The Structure of the Atom • Atoms are electrically neutral because they contain equal numbers of protons and electrons • The nuclei of atoms of different elements differ in their number of protons and therefore in the amount of positive charge they possess. • Thus, the number of protons determines that atom’s identity. Chapter 3 Section 2 The Structure of the Atom Properties of Subatomic Particles Chapter 3 Section 2 The Structure of the Atom Forces in the Nucleus When two protons are extremely close to each other, there is a strong attraction between them. A similar attraction exists when neutrons are very close to each other or when protons and neutrons are very close together. The short-range proton-neutron, proton-proton, and neutron-neutron forces that hold the nuclear particles together are referred to as nuclear forces. Chapter 3 Section 2 The Structure of the Atom The Sizes of Atoms • The radius of an atom is the distance from the center of the nucleus to the outer portion of its electron cloud. • Because atomic radii are so small, they are expressed using a unit that is more convenient for the sizes of atoms. • This unit is the picometer, pm. Objectives • Explain what isotopes are. • Define atomic number and mass number, and describe how they apply to isotopes. • Given the identity of a nuclide, determine its number of protons, neutrons, and electrons. • Define mole, Avogadro’s number, and molar mass, and state how all three are related. • Solve problems involving mass in grams, amount in moles, and number of atoms of an element. Atomic Number • Atoms of different elements have different numbers of protons. • Atoms of the same element all have the same number of protons. • The atomic number (Z) of an element is the number of protons of each atom of that element. Atomic Number Atomic Number Atomic number (Z) of element = # of p+ in nucleus of each atom of that element. Element Carbon # of protons Atomic # (Z) # of electrons 6 6 Phosphorus 15 15 Gold 79 79 electrically neutral atoms: # of p+ = # of e+ Isotopes and Mass Number • Isotopes are atoms of the same element that have different masses. • The isotopes of a particular element all have the same number of protons and electrons but different numbers of neutrons. • Most of the elements consist of mixtures of isotopes. • The mass number is the total number of protons and neutrons that make up the nucleus of an isotope. Mass Number Designating Isotopes • Hyphen notation: The mass number is written with a hyphen after the name of the element. • uranium-235 • Nuclear symbol: The superscript indicates the mass number and the subscript indicates the atomic number. 235 92 U • Nuclide is a general term for a specific isotope of an element. Mass Number Mass number is # of p+ and n0 in nucleus of an isotope: Mass # = p+ + n0 p+ n0 e- Mass # Oxygen - 18 8 10 8 18 Arsenic - 75 33 42 33 75 Phosphorus - 31 15 16 15 31 Isotope Sample Problem A How many protons, electrons, and neutrons are there in an atom of chlorine-37? Atomic number = 17 made up of 17 protons, electrons, 17. Mass number = 37 mass number − atomic number = 37 (p+ plus n0) − 17 p+ = 20 n0 An atom of chlorine-37 is made up of 17 e-, 17 p+, and 20 n0. Complete Symbols Superscript → Mass number Subscript → Atomic number Complete Symbols Find each of these: a) b) c) d) e) number of protons 35 number of neutrons 45 number of electrons 35 Atomic number 35 Mass Number 80 80 35 Br Complete Symbols If an element has an atomic number of 34 and a mass number of 79, what is the: a) b) c) d) number of protons 34 number of neutrons 45 number of electrons 34 complete symbol 79 34 Se Symbols If an element has 91 protons and 140 neutrons what is the a) Atomic number b) Mass number c) number of electrons d) complete symbol Symbols If an element has 78 electrons and 117 neutrons what is the a) Atomic number b) Mass number c) number of protons d) complete symbol Isotopes Dalton wrong about elements of same type being identical… Element’s atoms can have different # neutrons. different mass # Isotopes (flavors) Isotopes Frederick Soddy (1877-1956) proposed idea of isotopes in 1912 Isotopes - atoms of same element with different masses, b/c varying #s of neutrons 1921 Nobel Prize in Chemistry crater w/ his name on far side of Moon Naming Isotopes We can also put mass number after name of element: carbon-12 carbon-14 uranium-235 Isotope Protons Electrons Neutrons Hydrogen–1 (protium) 1 1 0 Hydrogen-2 (deuterium) 1 1 1 1 1 2 Hydrogen-3 (tritium) Nucleus Isotopes Elements occur in nature as mixtures of isotopes. Atomic Mass How heavy is an oxygen atom? Depends, b/c different kinds of oxygen atoms exist. We’re more concerned with average atomic mass. Based on abundance (%) of each variety of that element in nature. Don’t use grams - #’s tooooo small Measuring Atomic Mass Atomic Mass Unit (amu) one-twelfth mass of C-12 atom Carbon-12 Each chosen b/c of its isotope purity isotope has own atomic mass we determine average from % abundance. Calculate the average: Multiply atomic mass of each isotope by abundance (decimal), then add results. If not told otherwise, mass of isotope expressed in atomic mass units (amu) Atomic Masses Atomic mass is average of all naturally occurring isotopes of that element. Isotope Symbol Carbon-12 12C Carbon-13 13C Carbon-14 14C Composition of the nucleus 6 protons 6 neutrons 6 protons 7 neutrons 6 protons 8 neutrons % in nature 98.89% Atomic mass (amu) 12.00 1.11% 13.00 <0.01% 14.00 What is the average atomic mass of Carbon? 12.01