Download
energy and electrons n.
Skip this Video
Loading SlideShow in 5 Seconds..
Energy and Electrons PowerPoint Presentation
Download Presentation
Energy and Electrons

Energy and Electrons

2 Vues Download Presentation
Télécharger la présentation

Energy and Electrons

- - - - - - - - - - - - - - - - - - - - - - - - - - - E N D - - - - - - - - - - - - - - - - - - - - - - - - - - -
Presentation Transcript

  1. Energy and Electrons

  2. Electrons • What are electrons?

  3. Electron Configuration • Electron configuration is the arrangement of electrons in an atom, molecule or other body. • The electrons occupy specific probability regions (known as orbitals), whose shapes and electron capacity vary.

  4. Electron States • The lowest energy level that an electron can exist in is the ground state (n=1,2,…7). • Higher energy levels are called excited states (n>1). • In a given orbital, an electron that never radiates or absorbs energy is said to be in a stationary state.

  5. Energy as Light • If an atom absorbs exactly the right amount of energy, the electron will rise to the next orbital or energy level. • If an atom released an exact amount of energy, it would fall to a lower orbital or energy level. • The energy released appeared as a photon of light.

  6. Where are the electrons? • According to the Bohr model, electrons are found in energy levels around the nucleus. • There are 7 different energy levels • The farther the energy level is from the nucleus, the greater the amount of energy it holds.

  7. To Summarize ...

  8. For example ... • Calculate the number and types of sublevels, the total number of orbitals, and the maximum number of electrons with an energy level of 6 • n = 6 • types of orbitals = n = 6 • there are n2 orbitals on level 6 = 36 • there are 2n2 electrons = maximum 72 electrons

  9. Which state do you live in? • For each element, the atomic number = number of protons = number of electrons • The electron configuration must equal the number of protons in the atom and advance by 2n2 • Advance by 2 electrons, 4, 9, 16, etc.

  10. Electron Behavior • An electron can orbit in specified energy levels (orbitals). The further the orbital from the nucleus, the greater the energy level. • Orbital: a region in an atom where there is a high probability of finding electrons

  11. How many can fit? • Each level can hold a specific number of electrons • n = the energy level • 2n2 = the maximum number of electrons that energy level can hold • So, the first energy level can hold • 2 x 12 = 2 electrons

  12. Draw a Bohr Atom: Ground State Configuration • Put all of the protons in the nucleus • p+ • Determine the number of electrons in the atom • e- = p+ • Add your electrons to each energy level completely filling one level before moving to the next highest level. • Remember the max is 2n2 electrons in each level

  13. Bohr-ing Electrons! • Atomic Number = 13 13p+ 2e- 8e- 3e- For a total of 13 electrons

  14. You practice a few • Draw the ground-state electron configuration for • Oxygen • Potassium

  15. Oxygen • Atomic Number (# electrons) = 8 • Maximum # electrons = 8 = 2n2 • Level & Types of orbitals = n = 2 • Number of orbitals = n2 = 4 • Fill the electron orbitals like this: • 2 + 6

  16. Potassium • Atomic Number (# electrons) = 19 • n = 3.08 = 3 • Fill the orbitals 2 + 8 + 9

  17. Classwork • Draw the Bohr electron diagrams for the 1st 20 elements

  18. Quantum Numbers • To define the region in which electrons can be found, there are 4 quantum numbers assigned • Quantum number: a number that specifies the properties of electrons

  19. The principal of it all • The principal quantum number, n, indicates the main energy level occupied by the electron • The values for n are in positive whole integers (1, 2, 3, 4) • As n increases, the distance from the nucleus increases

  20. Next comes the l • The angular momentum quantum number, l, indicates the shape of the orbital • If l = 0, then there is an s orbital • If l = 1, p orbital • If l = 2, d orbital • If l = 3, f orbital

  21. Okay, now with the m • Next comes the magnetic quantum number, m, which is a subset of the l quantum number • This number indicates the numbers and orientations of the orbitals • The number of orbitals include 1s, 3p, 5d and 7f orbitals.

  22. Last, but not least,  • The spin quantum number, +1/2 or -1/2 (), indicates the orientation of the electron’s magnetic field relative to an outside magnetic field

  23. Orbitals • Each orbital is associated with a different letter: s, p, d, f, g, . . . • As chemists, we will only look at the s, p, d and f orbitals • Each orbital can accommodate only 2 electrons with opposite spins • Empty, half-filled and filled orbitals contain 0, 1 and 2 electrons, respectively

  24. Pauli Exclusion Principle • The principle that states that two particles of a certain class cannot be in the exact same energy state. • In other words, only 2 electrons can occupy a single orbital

  25. Heisenburg Uncertainty Principle • You cannot know simultaneously • Where an electron is, and • How fast it is moving

  26. S Orbitals • There is only one type of s orbital and it is present on every principal energy level. • The s orbital is spherical

  27. The p Orbital • There are three types of p orbitals (px, py and pz) • p orbitals are on every energy level except level 1 • All p orbitals have a dumbell shape

  28. The d Orbitals • There are 5 types of d orbitals and they are located on every energy level except for level 1 and level 2

  29. The f Orbitals • There are 7 types of f orbitals and they are located on every energy level except for levels 1, 2 and 3 • The shapes of the f orbitals are extremely complex

  30. Aufbau Principle • The principle that states that the structure of each successive element is obtained by adding one proton to the nucleus of the atom and one electron to the lowest energy orbital. • In other words, the electrons must fill the lowest energy level available

  31. The Diagonal Rule • The basic rule for assigning electrons to atoms is that electrons should occupy the lowest energy state possible. • To determine the relative energies or sublevels, use the diagonal rule. • Work from left and follow each arrow from tail to head and work from left to right.

  32. Start-up: October 7 • What is an electron configuration?

  33. The Chart 7s 7p 6s 6p 6d 5s 5p 5d 5f 4s 4p 4d 4f 3s 3p 3d 2s 2p 1s

  34. The Arrows 7s 7p 6s 6p 6d 5s 5p 5d 5f 4s 4p 4d 4f 3s 3p 3d 2s 2p 1s Start Here

  35. Write the electron filling pattern for the first 5sublevels. Element Atomic # Electron Config. H 1 1s1 He 2 1s2 Li 3 1s22s1 Be 4 1s22s2 B 5 1s22s22p1

  36. Shortcut • In order to conserve space, we use shorthand • When you reach a noble gas, the next element will begin a new principal energy level (He, Ne, Ar, Kr, Xe, Rn) • Place the symbol in brackets in place of the configuration scheme before it

  37. Sodium, for example • Neon Atomic Number = 10 • Electron Configuration: • 1s2 2s2 2p6 • Sodium Atomic Number = 11 • Electron Configuration: • [Ne] 3s1

  38. Classwork • Write the electron configuration of Bromine using the diagonal rule and the “shorthand” notation • 1s2 2s2 2p6 3s2 3p6 4s2 3d10 4p5 • [Ar] 4s2 3d10 4p5 or [Ar] 3d10 4s2 4p5

  39. Fun Facts • The orbital names s, p, d, and f stand for names given to groups of lines in the spectra of the alkali metals. These line groups are called sharp, principal, diffuse, and fundamental.

  40. Homework • P. 99 numbers 1 - 11