Nuclear Chemistry. Atomic Structure Notes Start on Slide 20 from the second class lecture

Transcription

1 Nuclear Chemistry Atomic Structure Notes Start on Slide 20 from the second class lecture

2 The Birth of an Idea Democritus, 400 B.C. coined the term atom If you divide matter into smaller and smaller pieces, eventually you will end up with tiny, indestructible pieces called atoms His ideas were rejected in favor of Aristotle s Aristotle 380 B.C.- all substances are made up of 4 elements Fire- hot Air- light Earth- cool, heavy Water- wet All substances a blend of these 4 elements

3 Dalton s Atomic Theory John Dalton, English teacher who did science on the side. Recorded his ideas in 3 points Dalton s Atomic Theory Dalton proposed that all matter is made up of tiny particles, which are molecules or atoms Molecules can be broken down into atoms by chemical processes Atoms cannot be broken down by chemical or physical processes

4 Dalton s Atomic Theory Law of Definite Composition: the % by mass of an element is always the same Ex: the mass ratio of carbon to oxygen in Carbon Dioxide (CO2) is always the same 1 carbon to 2 oxygen atoms Law of Conservation of Mass: In chemical reactions, mass is conserved and is not created nor destroyed Dalton proposed the creation of methane (CH4) by substituting 4 hydrogen atoms for the 2 oxygen atoms in carbon dioxide

5 Dalton s Atomic Theory 1. An element is composed of tiny, indestructible, indivisible particles called atoms 2. All atoms of the same element are identical, and have the same properties 3. Atoms of different elements combine to form compounds 4. Compounds contain atoms in small whole number ratios 5. Atoms can combine in more than one ratio to form different compounds, or simply, chemical reactions involve the rearrangement of atoms. No new atoms are created or destroyed

6 How small is an atom? An atom is the smallest particle of an element that retains the properties of that element. They are teeny tiny You would need to line up 100,000,000 Cu atoms to measure 1cm A penny that is made of pure Cu would contain 2.4 x 10^22 Cu atoms Atoms can be seen with a Scanning Tunneling Microscope (STM)

7 Structure of the Atom J.J. Thomson,1900- an English physicist proved that atoms had pieces called electrons (e-) Famous Cathode Ray Tube experiment Hooked up electrodes to a high-voltage source, created an anode (positively charged) and cathode (negatively charged) A glowing beam flowed from the negative disk to the positive disk, called the cathode ray The glowing beam was made up of negative particles (opposites attract)

8 Thomson s Plum Pudding Model 1. Atoms are breakable 2. e- are negative, so 1. Must be a positive charge to balance the e e- repel each other In the Plum Pudding Model: e- are suspended in a positively charged electric field Lots of empty space separates the e-

9 After Thomson Milliken (1900, American scientist) determined that the charge of an e- is -1; the mass is 9.11x10-28 g E. Goldstein discovered that a proton has a positive charge and is 1840 times heavier than the e- Proton s charge is +1; the mass is 1.67x10-24 g James Chadwick confirmed that the neutron has no charge, but the same mass as a proton Neutron s charge is 0; the mass is 1.67x10-24 g

10 Subatomic Particle Review Particle Symbol Relative Charge Relative Mass (mass of P+=1) Actual Mass (g) Electron e- 1-1/ x10-28 Proton p x10-24 Neutron N x10-24

11 Rutherford and Radioactivity There are 3 types of radiation: Alpha Particles (α) composed of positively charged helium nuclei Beta Particles (β) composed of negatively charged e- Gamma Particles (ϒ) composed of high energy radiation Size: α> β>ϒ

12 Rutherford s Gold Foil Experiment Ernest Rutherford- 1910, English physicist who believed in the plum pudding model Used radioactivity and shot the positively charged alpha particles at a gold foil which was a few atoms thick

13 Rutherford s Gold Foil Experiment Realized that the atom is largely empty space, so most of the α particles passed through the foil There is a dense, positive area at the center of the atom (named the nucleus) The α particles that deflected and bounced backwards did so after nearing or hitting the nucleus Rutherford predicted that because of the density of the nucleus, it must contain neutral particles in addition to protons 30 years later, Chadwick discovered neutrons

14 Bohr Model Niels Bohr, proposed Planetary Model Scientists realized that the attraction between e- and p+ should make the atom unstable Bohr proposed that e- occupy stable, fixed orbits around the nucleus with special quantized locations In the Bohr model, the e- can change orbits when it absorbs or gives off a photon of a specific color of light

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16 Quantum Theory Modern quantum theories lead to stable locations of e-, which are not exact orbits, but are characterized by specific quantum numbers e- may be found in clouds of probability, but the exact location of an e- can t be determined Every orbital has a different shape, and no 2 e- can be in the same orbital unless they have opposite spins (more on this later)

17 Review Questions Aristotle Suggested that matter existed through a combination of? Fire, Air, Earth, Water What were the 2 characteristics of the Dalton Model of the atom that were later found to be untrue? Atoms are indestructible/indivisible particles All atoms of the same element are identical What was most characteristic of Thomson s Model of the atom, that is most unlike what is accepted as true today? Plum pudding model, no nucleus, + charges mixed with e- How did Rutherford s experiments reveal that an atom consists of a dense, positively-charged nucleus, with a large e- cloud around it? Small ratio of α particles were deflected, and fewer bounced back after striking nucleus

18 Review Questions Why was the Bohr Model an improvement over the Rutherford Planetary model? Bohr proposed a model in which the e- would occupy stable, fixed orbits How did the Bohr Model explain the emission and absorption of light? e- can change orbits, accompanied by the absorption or emission of a photon of a specific color of light The Quantum Theory Model of the atom also describes Quantum levels, similar to Bohr. How did the new model differ? In quantum theory, the e- shells are not fixed orbits, but clouds of probability

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21 Atomic Notation Atomic Notation represents the number of protons and neutrons in the nucleus of an atom The Atomic Number, Z, represents the number of protons in the nucleus. It also represents the number of e- in a neutral atom. The Mass Number, A, represents the total number of protons and neutrons in the nucleus of an atom

22 Atomic Notation Atomic Number is 11 # of p+ = 22 protons; # e- = 11 electrons The Mass number is 23 #n= 23-11=12 neutrons

23 Atomic Notation Find the Number of protons Number of neutrons Number of electrons Atomic number Mass number

24 Atomic Notation 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?

25 Isotopes Atoms of the same element can have different numbers of neutrons (Dalton was wrong) These atoms of the same element would have different mass numbers These cousins of the same element are called isotopes Most elements occur naturally with varying number of neutrons

26 Hydrogen Isotopes Hydrogen has 3 Isotopes 1 proton and 0 neutrons, protium 1 proton and 1 neutron, deuterium 1 proton and 2 neutrons, tritium (radioactive)

27 Naming Isotopes To name an isotope properly, we put the mass number after the name of the element Carbon-12 (6 protons and 6 neutrons) Carbon-14 (6 protons and 8 neutrons) Uranium-235 (92 protons and 143 neutrons)

28 Practice Write the symbol for Cobalt-60 How many protons and neutrons does an atom of mercury-202 have? How many electrons are present in an atom of Copper-63?

29 Measuring Atomic Mass All of the elements on the Periodic Table have a Mass Number that is in decimal form, or not a whole number why? The Mass number reported on the periodic table is a weighted average of the isotopes that are naturally occurring

30 Measuring Atomic Mass The Mass Number= the Atomic Mass, or Atomic Weight The mass of 1 mole of an element (unit 4) The number of neutrons and protons present The mass of one atom of an element relative to an atom of another element Unit: Atomic Mass Unit (amu) The standard for the amu scale is Carbon: Carbon= 6p+ and 6n0= mass standard of 12 amu All other atoms are designated relative to Carbon. I.e. Hydrogen is 1/12 th the mass of carbon (or Atomic Mass 1 amu

31 Mathematics of Isotopes A weighted average means that the numbers of all the objects are not equal, but occur in different amounts Calculated by multiplying the percent of the object (as a decimal number) by its mass for each object and adding the numbers together

32 Mathematics of Isotopes Copper has 2 isotopes: Copper-63 with a mass of amu and 69.09% abundance Copper-65 with a mass of amu and 30.91% abundance amu amu = amu

33 Isotope Examples Magnesium has 3 isotopes % Magnesium-24 with a mass of amu, Magnesium-25 with a mass of amu, and the rest Magnesium-26 with a mass of amu. What is the atomic mass of Magnesium? Boron is 20% B-10 and 80% B-11. What is the Atomic Mass of Boron?

34 Ions Result from loss or gain of e- by an atom

35 Ions Cl

36 Ions What is the standard atomic notation for the aluminum atom and ion?