1 / 50

Atomic Theory

Atomic Theory. History of the Theory: Early ideas of the atom Democritus Aristotle John Dalton (pg. 56). Atomic Theory. History of the Theory: Early ideas of the atom Democritus - “atomos”; atoms exist Aristotle - atoms do NOT exist John Dalton (pg. 56). Atomic Theory.

marius
Download Presentation

Atomic Theory

An Image/Link below is provided (as is) to download presentation Download Policy: Content on the Website is provided to you AS IS for your information and personal use and may not be sold / licensed / shared on other websites without getting consent from its author. Content is provided to you AS IS for your information and personal use only. Download presentation by click this link. While downloading, if for some reason you are not able to download a presentation, the publisher may have deleted the file from their server. During download, if you can't get a presentation, the file might be deleted by the publisher.

E N D

Presentation Transcript

  1. Atomic Theory • History of the Theory: • Early ideas of the atom • Democritus • Aristotle • John Dalton (pg. 56)

  2. Atomic Theory • History of the Theory: • Early ideas of the atom • Democritus - “atomos”; atoms exist • Aristotle - atoms do NOT exist • John Dalton (pg. 56)

  3. Atomic Theory • History of the Theory: • Early ideas of the atom • Democritus - “atomos”; atoms exist • Aristotle - atoms do NOT exist • John Dalton (pg. 56) - 2000 years later; atoms exist • Matter is made up of atoms. • Atoms are indivisible. • Atoms of different elements are different. • LAW OF CONSERVATION OF MATTER

  4. More Recent Atomic Theory • E. Rutherford Nucleus in gold foil experiment • JJ Thomson • Electrons in cathode ray experiment

  5. Modern Atomic Theory • Neils Bohr Energy levels • Schrodinger • Sublevels & orbitals

  6. Subatomic Particles

  7. Proton Significance • Gives the atom its identity • Equivalent to the atomic number • Massive particle so adds to the atom’s mass • Gives the nucleus its positive charge • Balances the negative charge of the electrons

  8. Neutron Significance • Neutron glue that holds the nucleus together • Massive particle so adds to the atom’s mass (mass number= p + n) • Number can differ from one atom to another giving isotopes

  9. Electron Significance • Responsible for chemical properties • Forms ions when gained or lost, which leads to a charge

  10. Atomic Number • The picture is NOT completely correct. • Atomic # IS the number of protons in the atom. It is NOT always the number of electrons. • Therefore, the atomic # gives the atom its identity...its name.

  11. Atomic Number • In isotopic notation, the atomic number is shown in the lower left corner of the element symbol. • Example: 6C

  12. Atomic Number- Quick Check • If an atom is found to have 28 protons, what is its identity? • What is the atomic number of phosphorus? • How many protons does barium have? • Show the following elements in “isotopic notation”: • Lithium, sodium, sulfur, lead

  13. Charge • Charge = oxidation state • Oxidation state is just a fancier way to say it. • A charged atom results from an inequality between protons and electrons in an atom. • Which of those particles (protons or electrons) are more likely to be gained or lost from an atom? Why?

  14. Charge • Example: Calcium has 20 protons. It often loses electrons to other atoms. If calcium loses 2 electrons, the atom will have only 18 electrons. What will the charge be? • Answer: +2 charge

  15. Charge • In symbolic notation, the charge is listed in the top right corner of the element symbol. If the charge is neutral, then the corner is left blank. • Example: Cl-1 means a chlorine atom with 17 protons and 18 electrons. (It has one extra negative charge.)

  16. ChargeQuick Check • What is the charge of sulfur when it gains 2 electrons to the neutral atom? • Write the symbolic notation to represent an atom of aluminum that has lost 3 electrons leaving 13 protons and 10 electrons. • How many protons and electrons are in the following atom? Cu 2+

  17. Mass Number • Consider the subatomic particles. Which TWO particles have enough mass to matter? • Mass number is always a whole number...no decimals. • Mass number is calculated by adding the number of protons to the number of neutrons.

  18. Mass Number • In symbolic notation, the mass number is shown in the upper left corner of the element symbol. • Example: 126C

  19. Mass NumberQuick Check • If a sodium atom has 11 protons and 12 neutrons, what is the atom’s mass number? • How many neutrons does a copper atom with a mass number of 64 have? • How many protons and neutrons are in the atom represented with this symbolic notation: 3216S

  20. Isotopic Notation Summary • Shorthand way to record the element symbol, atomic number, charge, and mass number. • With this information, you can deduce the number of protons, neutrons, and electrons in the atom being represented. X Mass # Charge Atomic #

  21. Blocks on the Table • PROTONS give an atom its identity. They are the ONLY subatomic particle that must be identical from one atom to another. Neutrons AND electrons can vary... • SO, how can we create ONE box on the periodic table to represent ALL atoms of an element? We must represent AVERAGE atoms. (That’s why the atomic mass has decimals.)

  22. Atomic Mass • Atomic mass = the weighted average of all of the types of atoms of the element • Must know the % abundance and the mass number • Check out the example!

  23. Atomic Mass Example • Copper exists as a mixture of 2 types of atoms. The lighter copper has 29 protons and 34 neutrons, and it makes up 69.17% of all copper atoms. The heavier type has 29 protons and 36 neutrons. It makes up the remaining 30.83% of copper on earth. What is the atomic mass of copper?

  24. Atomic Mass Calculation • Chlorine has two isotopes. Chlorine-35 has an exact weight 34.968852 amu, and it has a 75.77% abundance. The other isotope has 36.965903 amu. What is the atomic mass?

  25. Isotopes TIME OUT! IF ALL THE ATOMS OF CARBON ARE NOT IDENTICAL, WHAT’S THE STORY WITH THE CARBON BLOCK ON THE PERIODIC TABLE? WHICH ONE DOES IT REPRESENT? • The three carbon atoms are ISOTOPES of carbon. • Definition: atoms of the same element that have different numbers of neutrons • Application: Does the difference of electrons matter when considering ISOTOPES? • Since neutrons are massive, a change in the number of neutrons gives a different mass number. • Mass number = protons + neutrons • “Sisters” - all with the same # of protons but slightly different masses

  26. Isotope/Atomic Mass Lab • Try it yourself now! • Get your sample of Candium. • Draw the data table on your notebook paper. • Gather the data, and calculate the atomic mass of candium. • You can eat your sample when you’re done! 8-)

  27. Energetic Electron • Electrons are energetic, and they exist on energy levels. • Quantum: specific amount of energy needed to move from one energy level to the next; energy levels are given principal quantum numbers as a result • During this movement, energy is absorbed and released. • Light is sometimes visible when the energy is released.

  28. Quantum Illustration Quantum = distance between two energy levels Principal Quantum Number (n) = begin counting closest to the nucleus

  29. Quantum Leap Illustration An electron from n=2 (ground state) can absorb a quantum of energy and jump to n=3 (excited state). Excited state is temporary. The electron will soon release the quantum and fall back to ground state. The released energy will travel in a wave.

  30. Electromagnetic Spectrum(All energy waves fall into one of these categories.)

  31. Calculating the Emission of Energy • Two Equations to calculate the energy: • Speed of light = wavelength x frequency c = λν c = 3.0 x108 m/s (speed of light is constant) 2. Energy = Planck’s constant x frequency E = hν h=6.626x10-34 Jsec (Planck’s constant)

  32. iRespond Question Multiple Choice F Finding a Connection Look at the two equations. What variable do they have in common? A.) Speed of light (c) B.) Wavelength (lambda) C.) Frequency (nu) D.) Energy (E) E.)

  33. Electromagnetic Wave Calculations • So, if I know the _______ of a wave, then I can calculate its wavelength AND its energy.

  34. Energy Practice Problem: If the wavelength is 0.001 meters, then what is the frequency of the wave? iRespond Question Multiple Choice F A.) 300,000 Hz B.) 3.0x1011 Hz C.) 3.33x10-12 Hz D.) 6.626x10-37 Hz E.)

  35. Energy Practice Problem: If the wavelength is 0.001 meters, then what is the energy of the wave? iRespond Question Multiple Choice F A.) 1.99x10-22 J B.) 2.21x10-45 J C.) 1.99x1046 J D.) 6.626x10-37 J E.)

  36. iRespond Question Multiple Choice F Energy Practice Problem: If the frequency is 7.0x1013 Hz, what is the wavelength of the wave? What is the energy of the wave? A.) 4.29x106 m; 4.63x1020 J B.) 2.33x105 m; 1.55x10-28 J C.) 4.29x10-6 m; 4.63x10-20 J D.) None of these E.)

  37. Periodic Table Labeling S block: blue d block: red P block: yellow f block: green

  38. Hotel Tarvin • Managers: • Aufbau: Each room/bed on the lower floors must be occupied before moving to a higher floor. • Pauli: A maximum of two guests may occupy a bed. Guests must sleep head to foot. • Hund: Single guests must occupy separate beds in a room. No pairing occurs unless no empty beds exist.

  39. Hotel Tarvin Rooms available: S: Single room, has only one bed (sleeps 2) P: Prestige room, has three beds (sleeps 6) D: Deluxe room, has five beds (sleeps 10) F: Fabulous room, has 7 beds (sleeps 14)

  40. The analogy... • Hotel = electron cloud • Floors = energy levels • 1-7 COEFFICIENTS • Rooms = sublevels (shapes) • s = sphere, p = dumbbell, d & f • Beds = orbitals ( & axis) • x, y,and z • Guests = electrons • Up and down = direction of spin

  41. Let’s practice... • Write the electron configuration for: • Fluorine • Magnesium • Bromine • Palladium • Xenon • Write the noble gas notation for the elements above.

  42. What does this all mean? • Let’s analyze the configuration of magnesium. • Energy levels • Sublevels • Electrons • Valence level and electrons • What can I NOT find represented in the electron configuration? • Orbitals • Pairing

  43. Orbital Diagrams • Visual representation of electron cloud • Shows everything that electron configuration shows PLUS orbitals & pairing • How to Draw an Orbital Diagram • Write the electron configuration first. • Draw the “beds” (orbitals) under each room. • Put the electrons in the “beds.” • Let’s practice with the electron configurations written already.

  44. Analyzing Orbital Diagrams & Stability of Atoms • Electrons are responsible for the chemical properties (personalities) of atoms. • STABLE = All orbitals (beds) are full. • PRETTY STABLE = All orbitals (beds) are full OR half-full. • UNSTABLE = Empty orbitals (beds) exist. NOTE: If an atom has to have an empty orbital, it is best for the empty orbital to be on the valence energy level so that it might be filled by electrons from a nearby atom.

  45. iRespond Question Multiple Choice F Analyzing Stability Draw the orbital diagram of sodium. Is sodium stable, pretty stable, or unstable? A.) Stable B.) Pretty Stable D.) C.) Unstable E.)

  46. Draw the orbital diagram of chromium. Is Cr stable, pretty stable, or unstable? iRespond Question Multiple Choice F Analyzing Stability A.) Stable B.) Pretty Stable C.) Unstable D.) E.)

  47. Practice Analyzing Stability 1. Draw the orbital diagram of sodium. Is sodium stable, pretty stable, or unstable? • Answer: UNSTABLE; Sodium has one half-full orbital, but it’s third energy level has three empty p orbitals. • Draw the orbital diagram of chromium. Is Cr stable, pretty stable, or unstable? • Answer: UNSTABLE; chromium has one empty orbital.

  48. Exceptions to the Rules • Exceptions to Aufbau principle: • Chromium and Copper columns ONLY • Stability of filled and half-filled sublevels • Arrangement’s impact on bonding

  49. Atomic Emission Spectra • Unique to each element • “Fingerprint” • Used to identify unknowns • Shows all wavelengths of visible light given off by an atom • Flame Test Lab

More Related