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Something Smaller Than An Atom?

Something Smaller Than An Atom?. Atomic Structure. Review of the Atom. Smallest part of matter representing an element Once was thought to be the smallest part of matter Later, scientists discovered atoms are made of subatomic particles. Subatomic particles. Protons - positive charge

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Something Smaller Than An Atom?

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  1. Something Smaller Than An Atom? Atomic Structure

  2. Review of the Atom • Smallest part of matter representing an element • Once was thought to be the smallest part of matter • Later, scientists discovered atoms are made of subatomic particles

  3. Subatomic particles • Protons - positive charge • Neutrons – no charge (neutral) • Electrons – negative charge

  4. Representing a Specific Element A X Z Mass Number (P + + No) Atomic Number (P+) Atomic Symbol

  5. Nucleus: Center Stage • Ernest Rutherford discovered atoms have a nucleus (1911) Later scientists discovered that the nucleus also contains all the neutrons

  6. Rutherford’s Model • Discovered the nucleus • Electrons moved around in Electron cloud • Mostly empty space

  7. Where are those electrons? • Early models of the atom showed electrons spinning around the nucleus randomly • Research showed that this is NOT true

  8. Bohr’s Model • Why don’t the electrons fall into the nucleus? • Move like planets around the sun. • In circular orbits at different levels. • Energy separates one level from another.

  9. Bohr’s Model Nucleus Electron Orbit Energy Levels

  10. More about Bohr….. • Created a model showing electrons are in orbits of different energy around the nucleus • …… Think of the planets orbiting the sun Bohr used the term energy levels (or shells) to describe these orbits with different amounts of energy. He also said that the energy of an electron is quantized, meaning electrons can have one energy level or another but nothing in between.

  11. Bohr’s Model } Fifth • Further away from the nucleus means more energy. • There is no “in between” energy • Energy Levels Fourth Third Increasing energy Second First Nucleus

  12. Let’s get close to the nucleus • Bohr found that the closer an electron is to the nucleus, the less energy it needs. • The farther away it is, the more energy it needs • He numbered the electron’s energy levels

  13. Energy levels • 1st level can hold up to 2 electrons • 2nd level can hold up to 8 electrons • 3rd level can hold up to 18 electrons • And so on……..

  14. Energy levels • E-level an electron normally occupies is called ground state • But, electrons can move to a higher –energy, less-stable level, or shell, by absorbing energy. This higher energy, less-stable state is called the electron’s excited state.

  15. The electron gets crunk….. • When the electron is finished being excited it goes back to its ground state by releasing some of the energy it has absorbed

  16. Line spectrum…..what is that??? • Energy released by electrons sometimes occupies part of the electromagnetic spectrum that humans detect as visible light

  17. Problem with Bohr’s model • Unexplainable observations on complex atoms until the quantum theory was created Quantum theory ----- matter has properties associated with waves. It is impossible to know the exact position and momentum (speed and direction) of an electron at the same time. ------UNCERTAINTY PRINCIPLE

  18. The Quantum Mechanical Model • Energy is quantized. It comes in chunks. • Quanta - the amount of energy needed to move from one energy level to another. • Quantum leap in energy. • Treated electrons as waves

  19. The Quantum Mechanical Model • Does have energy levels for electrons. • Orbits are not circular. • It can only tell us the probability of finding an electron a certain distance from the nucleus.

  20. The Quantum Mechanical Model • The electron is found inside a blurry “electron cloud” • A area where there is a chance of finding an electron.

  21. Quantum mechanical model of the atom Pay attention so you don’t get LOST !!!!!!!

  22. Atomic Orbitals • Principal Quantum Number (n) = the energy level of the electron. • Within each energy level the complex math of Schrödinger's equation describes several shapes. • These are called atomic orbitals • Regions where there is a high probability of finding an electron.

  23. S orbitals • 1 s orbital for every energy level • Spherical shaped • Each s orbital can hold 2 electrons • Called the 1s, 2s, 3s, etc.. orbitals.

  24. P orbitals • Start at the second energy level • 3 different directions • 3 different shapes (dumbell) • Each can hold 2 electrons

  25. D orbitals • Start at the third energy level • 5 different shapes • Each can hold 2 electrons

  26. F orbitals • Start at the fourth energy level • Have seven different shapes • 2 electrons per shape

  27. Summary # of shapes Max electrons Starts at energy level s 1 2 1 p 3 6 2 5 10 3 d 7 14 4 f

  28. Bed check for Electrons…Electron Configurations…… • Goal: Use an energy level diagram to depict electrons for any element • 1s orbital is closest to the nucleus and it has the lowest energy • At energy level 2, there are both s and p orbitals, with the 2s having lower energy than the 2p.

  29. Energy level diagram continued • The three 2p subshells are represented by three dashes of the same energy. • Energy levels 3, 4, and 5 are also shown • Notice, that 4s has lower energy than the 3d. This is an exception to what you thought but it does occur in nature like this.

  30. 7p 6d 5f 7s 6p 5d 6s 4f 5p 4d 5s 4p 3d 4s 3p Increasing energy 3s 2p 2s 1s

  31. Electron Configurations • The way electrons are arranged in atoms. • Aufbau principle- electrons enter the lowest energy first. • This causes difficulties because of the overlap of orbitals of different energies. • Pauli Exclusion Principle- at most 2 electrons per orbital - different spins

  32. Aufbau Principle • Method for remembering the order in which orbitals fill the vacant energy levels (like a people fill up vacant rooms in a hotel) RULE: ELECTRONS FILL THE LOWEST VACANY ENERGY LEVELS FIRST. ANOTHER RULE: WHEN THERE’S MORE THAN ONE SUB-SHELL AT A PARTICULAR ENERGY LEVEL, SUCH AS AT THE 3P OR 4 D LEVELS, ONLY ONE ELECTRON FILLS EACH SUB-SHELL UNTIL EACH SUBSHELL HAS ONE ELECTRON. THEN, ELECTRONS START PAIRING UP IN EACH SUBSHELL. THIS RULE IS CALLED HUND’S RULE.

  33. Electron Configuration • Hund’s Rule- When electrons occupy orbitals of equal energy they don’t pair up until they have to . • Let’s determine the electron configuration for Phosphorus • Need to account for 15 electrons

  34. 7p 6d 5f 7s 6p 5d 6s 4f 5p 4d 5s 4p 3d 4s 3p Increasing energy 3s 2p 2s 1s • The first to electrons go into the 1s orbital • Notice the opposite spins • only 13 more

  35. 7p 6d 5f 7s 6p 5d 6s 4f 5p 4d 5s 4p 3d 4s 3p Increasing energy 3s 2p 2s 1s • The next electrons go into the 2s orbital • only 11 more

  36. 7p 6d 5f 7s 6p 5d 6s 4f 5p 4d 5s 4p 3d 4s 3p Increasing energy 3s 2p 2s 1s • The next electrons go into the 2p orbital • only 5 more

  37. 7p 6d 5f 7s 6p 5d 6s 4f 5p 4d 5s 4p 3d 4s 3p Increasing energy 3s 2p 2s 1s • The next electrons go into the 3s orbital • only 3 more

  38. 7p 6d 5f 7s 6p 5d 6s 4f 5p 4d 5s 4p 3d 4s 3p Increasing energy 3s 2p 2s 1s • The last three electrons go into the 3p orbitals. • They each go into separate shapes • 3 unpaired electrons • 1s22s22p63s23p3

  39. 7s 7p 7d 7f 6s 6p 6d 6f 5s 5p 5d 5f 4s 4p 4d 4f 3s 3p 3d 2s 2p 1s The easy way to remember • 1s2 • 2 electrons

  40. 7s 7p 7d 7f 6s 6p 6d 6f 5s 5p 5d 5f 4s 4p 4d 4f 3s 3p 3d 2s 2p 1s Fill from the bottom up following the arrows • 1s2 2s2 • 4 electrons

  41. 7s 7p 7d 7f 6s 6p 6d 6f 5s 5p 5d 5f 4s 4p 4d 4f 3s 3p 3d 2s 2p 1s Fill from the bottom up following the arrows • 1s2 2s2 2p6 3s2 • 12 electrons

  42. 7s 7p 7d 7f 6s 6p 6d 6f 5s 5p 5d 5f 4s 4p 4d 4f 3s 3p 3d 2s 2p 1s Fill from the bottom up following the arrows • 1s2 2s2 2p6 3s2 3p6 4s2 • 20 electrons

  43. Draw the energy level diagram of oxygen • Find oxygen on the period table • Atomic number = 8 • Therefore, it has 8 protons and 8 electrons • So, you put 8 electrons into your energy level diagrams

  44. Energy level diagram for oxygen • Use arrows to represent electrons • If two electrons fit in the same orbital, one must face up and the other one down The first electron goes into the 1s orbital, filling the lowest energy level first. And, the second one spin pairs with the first one. Electrons 3 and 4 spin pair in the next lowest vacant orbital – the 2 s. Electron 5 goes into one of the 2p sub-shells Electrons 6 and 7 go into the other two totally vacant 2p orbitals The last electron spin pairs with one of the electrons in the 2p subshells.

  45. Electron Configuration • Oxygen 1s22s22p4 • You can derive the electron configuration from the energy level diagram. • The first two electrons in oxygen fill the 1s orbital, so you it as 1s2 in the electron configuration. • The 1 is the energy level, the s represents the type of orbital, and the superscript 2 represents the number of electrons in that orbital. • The next two electrons are in the 2s orbital, so you write 2s2. • Last, you show the 4 electrons in the 2p orbital as 2p4. • That’s how you get 1s22s22p4.

  46. Tip • The sum of the superscript numbers equals the atomic number, or the number of electrons in the atom.

  47. Exceptions to Electron Configuration

  48. Orbitals fill in order • Lowest energy to higher energy. • Adding electrons can change the energy of the orbital. • Filled and half-filled orbitals have a lower energy. • Makes them more stable. • Changes the filling order of d orbitals

  49. Copper’s electron configuration • Copper has 29 electrons so we expect • 1s22s22p63s23p64s23d9 • But the actual configuration is • 1s22s22p63s23p64s13d10 • This gives one filled orbital and one half filled orbital. • Remember these exceptions • s2d4 s1d5 • s2d9 s1d10

  50. Here are a couple of electron configurations you can use to check your conversions from energy level diagrams: Chlorine (Cl) Iron (Fe)

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