Atomic Structure Atoms and their structure Mr. Bruder

Democritus (400 B.C.) proposed that matter was composed of tiny indivisible particles called atomos not scientifically tested "Nothing exists but atoms and empty space; everything else is opinion."

Democritus (400 B.C.) rejected by Aristotle and others who believed that matter could be endlessly divided

2-4 Evidence for Atoms n Law of Conservation of Mass –Mass is not gained or lost in a chemical reaction. n Proposed by Antoine Lavoisier in n What would happen to the mass reading if the reaction was done without the balloon (an open system)? Figure 2.2

2-5 Evidence for Atoms n Law of Definite Proportions –Proposed by Joseph Proust between 1797 and 1804 –A compound always has the same relative amounts of the elements that compose it. –For example, when water is broken down by electrolysis into oxygen and hydrogen, the mass ratio is always 8 to 1. Figure 1.2

John Dalton (1807) British Schoolteacher n based his theory on others’ experimental data and developed the atomic theory Billiard Ball Model n atom is a uniform, solid sphere

Dalton’s Atomic Theory n John Dalton ( ) had four theories 1. All elements are composed of submicroscopic indivisible particles called atoms 2. Atoms of the same element are identical. The atoms of anyone element are different from those of any other element 3. Atoms of different elements can physically mix together or can chemically combine w/ one another in simple whole-number ratios to form compounds 4. Chemical reactions occur when atoms are separated, joined, or rearranged. However, atoms of one element are never changed into atoms of another elements as a result of a chemical reaction

Atoms & Subatomic Particles n Atom- smallest particle of an element that retains the properties of that element

2-9 Proof of Atoms: STM Image of Gold n The scanning tunneling microscope, invented in 1981, allows us to create images of matter at the atomic level. Figure 2.4

Electron n J.J Thomson ( ) – discovered the electron in 1897 n Electron is the negative charged subatomic particle n An electron carries exactly one unit of negative charge & its mass is 1/1840 the mass of a hydrogen atom

Cathode Ray n The Cathode Ray tubes pass electricity through a gas that is contained at a very low pressure

Thomson’s Experiment Voltage source +-

Thomson’s Experiment Voltage source +-

Thomson’s Experiment Voltage source +-

n Passing an electric current makes a beam appear to move from the negative to the positive end Thomson’s Experiment Voltage source +-

n Passing an electric current makes a beam appear to move from the negative to the positive end Thomson’s Experiment Voltage source +-

n Passing an electric current makes a beam appear to move from the negative to the positive end Thomson’s Experiment Voltage source +-

n Passing an electric current makes a beam appear to move from the negative to the positive end Thomson’s Experiment Voltage source +-

Thomson’s Experiment n By adding an electric field

Voltage source Thomson’s Experiment n By adding an electric field + -

Voltage source Thomson’s Experiment n By adding an electric field + -

Voltage source Thomson’s Experiment n By adding an electric field + -

Voltage source Thomson’s Experiment n By adding an electric field + -

Voltage source Thomson’s Experiment n By adding an electric field + -

Voltage source Thomson’s Experiment n By adding an electric field he found that the moving pieces were negative + -

Thomson’s Atomic Model n Thomson’s Atomic Model n Thomson though electrons were like plums embedded in a positively charged “pudding”, so his model was called the “plum pudding” model

Thomsom’s Model n Found the electron n Couldn’t find positive (for a while) n Said the atom was like plum pudding n A bunch of positive stuff, with the electrons able to be removed

Mass of Electron n In 1909 Robert Millikan determined the mass of an electron with his Oil Drop Experiment n He determined the mass to be x kg n The oil drop apparatus

Millikan’s Experiment Atomizer Microscope - + Oil

Millikan’s Experiment Oil Atomizer Microscope - + Oil droplets

Millikan’s Experiment X-rays X-rays give some drops a charge by knocking off electrons

Millikan’s Experiment +

They put an electric charge on the plates

Millikan’s Experiment Some drops would hover

Millikan’s Experiment

Measure the drop and find volume from 4/3πr 3 Find mass from M = D x V

Millikan’s Experiment From the mass of the drop and the charge on the plates, he calculated the charge on an electron

Proton n In 1886 Goldstein discovered the Proton n Proton is a positively charged subatomic particle found in the nucleus of a atom

Radioactivity n Discovered by accident n Bequerel n Three types –alpha- helium nucleus (+2 charge, large mass) –beta- high speed electron –gamma- high energy light

Ernest Rutherford n Rutherford ( ) proposed that all mass and all positive charges are in a small concentrated region at the center of the atom n He used the Gold-Foil Experiment to prove his theory n In 1911 he discovered the Nucleus n Nucleus- central core of an atom, composed of protons and neutrons n The nucleus is a positively charged region and it is surrounded by electrons which occupy most of the volume of the atom

Rutherford’s Experiment n Used uranium to produce alpha particles n Aimed alpha particles at gold foil by drilling hole in lead block n Since the mass is evenly distributed in gold atoms alpha particles should go straight through. n Used gold foil because it could be made atoms thin

Lead block Uranium Gold Foil Florescent Screen

What he expected

Because

Because, he thought the mass was evenly distributed in the atom

What he got

How he explained it + n Atom is mostly empty n Small dense, positive piece at center n Alpha particles are deflected by it if they get close enough

+

Copyright © Cengage Learning. All rights reserved49 Nuclear Atom Viewed in Cross Section

Neutron n James Chadwick ( ) – discovered the neutron in 1932 n Neutron is a subatomic particle with no charge but their mass is nearly equal to that of a proton

2-51 The Neutron n Because the protons in the atom could account for only about half the mass of most atoms, scientists knew there was another heavy particle in the nucleus. n Neutrons were proposed by Ernest Rutherford in 1907 (to account for a mass discrepancy in the nucleus) and discovered in 1932 by James Chadwick. n The neutron has about the same mass as a proton but with no charge.

James Chadwick (1932) Discovered neutrons n neutral particles in the nucleus of an atom Joliot-Curie Experiment n based his theory on their experimental evidence

James Chadwick (1932) Joliot-Curie Experiment Be Gamma rays? 50 MeV Energy values weren’t consistent. 5 MeV Alpha Particles

James Chadwick (1932) Chadwick’s Explanation Be 5 MeV Alpha Particles Energy values were consistent. Also accounted for extra mass in the nucleus. Neutrons 5 MeV

James Chadwick (1932) Neutron Model

56 Bohr Model n Bohr changed the Rutherford model and explained how the electrons travel. n Bohr explained the following in his model: 1. Electrons travel in definite orbits with a certain energy around the nucleus. They must gain or lose energy in certain packages called Quanta 2. He explained that accelerating particles should get pulled into the nucleus but that does not occur because of the stability of the atom 3. His model was patterned after the motion of the planets around the sun. It is often called the Planetary model.

Bohr’s Model Nucleus Electron Orbit Energy Levels

The Quantum Mechanical Model n A totally new approach n Several different people made important contributions to the development of the model.

De Broglie n De Broglie said matter could be like a wave. n De Broglie said particles like the electron could now how properties of both a waves and particles

Heisenberg Uncertainty Principle n It is impossible to know exactly the position and velocity (momentum) of a particle. n The better we know one, the less we know the other. n The act of measuring changes the properties. n More precisely the velocity is measured, less precise is the position (vice versa).

n Things that are very small behave differently from things big enough to see. n The quantum mechanical model is a mathematical solution. He applied these solutions to waves of electrons n It is not like anything you can see. The Quantum Mechanical Model Schrodinger

Erwin Schrödinger (1926) Quantum mechanics n electrons can only exist in specified energy states Electron cloud model n orbital: region around the nucleus where e - are likely to be found

Erwin Schrödinger (1926) Electron Cloud Model (orbital) dots represent probability of finding an e -

Modern View n The atom is mostly empty space n Two regions n Nucleus- protons and neutrons n Electron cloud- region where you have a chance of finding an electron

Quark n Protons & Neutrons can still be broken down into a smaller particle called the Quark n The Quark is held together by Gluons

Density and the Atom n Since most of the particles went through, it was mostly empty. n Because the pieces turned so much, the positive pieces were heavy. n Small volume, big mass, big density n This small dense positive area is the nucleus

Chapter 2: Atoms, Molecules, and Ions67 Subatomic Particles Protons and neutrons are located at the center of an atom called the nucleus. Electrons are dispersed around the nucleus. EOS

Atomic Particles ParticleChargeMass (kg)Location Electron9.109 x Electron cloud Proton x Nucleus Neutron x Nucleus

Subatomic particles Electron Proton Neutron NameSymbolCharge Relative mass Actual mass (g) e-e- p+p+ n0n / x x

Symbols n Contain the symbol of the element, the mass number and the atomic number X Mass number Atomic number

Sub-atomic Particles n Z - atomic number = number of protons determines type of atom n A - mass number = number of protons + neutrons n Number of protons = number of electrons if neutral

Symbols X A Z Na 23 11

Atomic Structure Symbols n Proton = p + n Electron = e - n Neutron = n 0 n Atomic # - Subscript n Mass # - Superscript

Rules for Atomic Structure 1. Atomic # = # of Protons 2. # of Protons = # of Electrons 3. Mass # = # of Protons + # of Neutrons n # of Neutrons = Mass # - # of Protons n If you know the Mass # & Atomic # you know the composition of the element

Symbols n Find n Find the –number –number of protons of neutrons of electrons –Atomic –Atomic number –Mass –Mass Number Br 80 35

Symbols n if an element has an atomic number of 34 and a mass number of 78 what is the –number of protons –number of neutrons –number of electrons –Complete symbol

Symbols n if an element has 78 electrons and 117 neutrons what is the –Atomic number –Mass number –number of protons –Complete symbol

Example Element Atomic # Mass #ProtonsElectro ns Neutro ns K

Isotopes n Isotope- atoms that have the same number of protons but different number of neutrons n Since isotopes have a different number of neutrons the isotope has a different mass number. n Isotopes are still chemically alike because they have the same number of protons and electrons

Copyright © Cengage Learning. All rights reserved80 Two Isotopes of Sodium

Examples of Isotopes

2-82 Heavy Water n One ice cube is made with water that contains only the hydrogen-2 isotope. The other ice cube is composed of water with normal water which contains mostly hydrogen-1. n Which is which? Figure 2.13

© 2009, Prentice-Hall, Inc. Isotopes n Isotopes are atoms of the same element with different masses. n Isotopes have different numbers of neutrons C 12 6 C 13 6 C 14 6 C

Naming Isotopes n Put the mass number after the name of the element n carbon- 12 n carbon -14 n uranium-235

Electrical Charges n Electrical charges are carried by particles of matter n Atoms have no net electrical charges n Given the number of negative charges combines with the number of positive charges = Electrically Neutral n All elements are Electrically Neutral

Atomic Mass vs. Atomic Weight n Atomic Mass is for a single element n Most elements are Isotopes n How do we find their mass? n We use Atomic Weight

Chapter 2: Atoms, Molecules, and Ions87 Atomic Masses EOS An atomic mass unit (amu) is defined as exactly one-twelfth the mass of a carbon-12 atom 1 u = × 10 –24 g The atomic mass of an element is the weighted average of the masses of the naturally occurring isotopes of that element

Measuring Atomic Mass n Unit is the Atomic Mass Unit (amu) n One twelfth the mass of a carbon-12 atom n Each isotope has its own atomic mass. We need the average from the percent abundance n Each isotope of an element has fixed mass and a natural % abundance n You need both of these values to find the Atomic Weight

Calculating Atomic Weight n Cl amu and 75.77% abundance n Cl amu and 24.23% abundance n To solve for Cl AMU x Abundance x = n You solve for Cl-37

Atomic Weight Cont. n Cl AMU x Abundance x = n Now you combine your two answers n = n n Look at Cl on the table. What is the Atomic Weight?

Example n Calculate the atomic weight of copper. Copper has two isotopes. One has 69.1% and has a mass of amu. The other has a mass of amu. What is the atomic weight???

Atomic Weight & Decimals n Atomic Weight- of an element is a weighted average mass of the atoms in a naturally occurring sample of an element n Atomic Weights use decimal points because it is an average of an element