1 / 25

Electrons

Electrons. Review of Atoms and Atomic Theory. So far we have described the atom as a nucleus of protons and neutrons surrounded by electrons – this works well for a simple explanation, but does not explain certain properties of elements.

vielka-lang
Télécharger la présentation

Electrons

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. Electrons

  2. Review of Atoms and Atomic Theory • So far we have described the atom as a nucleus of protons and neutrons surrounded by electrons – this works well for a simple explanation, but does not explain certain properties of elements. • For example, why do metals give off a characteristic color when heated in a flame? Answer: Because the electrons are giving off light energy as they return to a lower energy level. More about this later!!!

  3. Dalton’s Atomic Theory, Review: • All elements are composed of tiny indivisible particles called atoms • Atoms of the same element are identical. The atoms of one element are different from atoms of a different element • Atoms of different elements can physically mix together, or chemically combine in simple whole-number ratios to form compounds • Chemical reactions occur when atoms are separated, joined or rearranged.

  4. JJ Thompson • JJ Thompson (who discovered the electron) realized that Dalton’s model did not take into account electrons – so he proposed a revised model often called the plum-pudding model • This had a central positively charged mass onto which electrons were just stuck to the outside. • It did not describe the number of protons and electrons or how they were arranged.

  5. Ernest Rutherford • After discovering the nucleus, Ernest Rutherford proposed the nuclear atom, in which electrons surrounded a dense nucleus – leaving lots of empty space • Later experiments showed a nucleus of + charged protons and neutrally charged neutrons, surrounded by – charged electrons. • So why didn’t the – charged electrons fall into the + charged protons of the nucleus?

  6. Bohr Model of Electrons • In the early 1900’s, Niels Bohr, who was a student of Rutherford, came up with a new atomic theory • Bohr proposed that electrons were in circular orbits around the nucleus, much like planets orbit the sun (thus it was called the planetary model) • Bohr suggested that electrons in a particular orbit have a fixed energy, and that electrons do not lose energy and cannot fall into the nucleus.

  7. Energy Levels of Electrons • The energy level of an electron is the region around the nucleus where the electron is likely to be moving. These are also called electron shells. • These energy levels are like the rungs on a ladder – just like you can not stand between rungs on a ladder, electrons cannot be between energy levels.

  8. Quantums of Energy • A quantum of energy is the amount of energy needed to move an electron from its present energy level to the next high level. • The term quantum leap comes from describing this sudden change from one level to the next. • The higher the energy level, the farther the electron typically is from the nucleus.

  9. The Quantum Mechanical Model (The Cloud Model) • Finally, in 1926 Erwin Schrödinger took the atom to its current shape • The modern description of the electrons in an atom is called the quantum mechanical model and is based on the mathematical equations from Schrödinger.

  10. The Cloud Model • The Quantum Mechanical Model (also called the Cloud Model) does not describe electrons in an orbit like the planets. • The probability of finding an electron within a certain volume of space surrounding the nucleus is represented as fuzzy cloud, which represents where an electron would be found 90% of the time.

  11. Atomic Orbitals • In the Bohr model we had orbits, but since the quantum mechanical model does not have orbits, but rather regions of probability, the regions are called atomic orbitals. (They may also be called electron clouds.)

  12. Principle Quantum Number-the main Energy Level occupied by an electron • The main energy level of an electron (or shell) is designated by means of a principle quantum number (n) • This is assigned in order of increasing energy levels: (shells) n =1,2,3,4,5,. . . • Within each shell, there are sublevels. The number of sublevels found will be the same as the value of n.

  13. Where will the electrons be found? • In each shell, electrons will be found in regions of probability called atomic orbitals. • Each orbital is given a letter symbol: S, P, D, or F: primarily based on shape • Each orbital can hold two electrons.

  14. S, P, D, or F: Shapes of Orbitals • S orbitals are spherical • P orbitals are shaped like dumbbells (or peanuts): there are 3 of these: x, y, and z • D orbitals are like two sets of dumbbells at right angles to each other • F orbitals are too complex to visualize

  15. Sublevel 1: 1s • For the first shell (n =1), there is one sublevel called 1s. • In the s atomic orbital, there is an equal probability of finding an electron in any direction from the nucleus S=spherical

  16. Sublevel 2: 2s and 2p • For the second shell (n =2), there are two sublevels: 2s and 2p • The 2s orbital is spherical, while the 2p orbitals are dumbbell shape • The 2p orbital is of higher energy level than the 2s orbital • The 2p orbital consists of three p orbitals, each aligned on the x, y & z axis • Thus the second energy level has four orbitals 2s, 2z, 2y, 2z

  17. Sublevel 3: 3s, 3p, and 3d • For the third shell (n=3), there are 3 sublevels: 3s, 3p and 3d • As before, there is one 3s orbital and three 3p orbitals – but 3d has 5 orbitals, for a total of 9 orbitals • How many electrons can Sublevel 3 hold?

  18. Sublevel 4: 4s, 4p, 4d, and 4f • For the fourth shell (n=4), there are 4 sublevels: 4s, 4p, 4d and 4f • The 4th principle energy level has 16 orbitals: One 4s, three 4 p, five 4d and seven 4f orbitals • How many electrons can the 4th sublevel hold?

  19. Electron Configuration • The arrangement of electrons in an atom is usually shown by writing an electron configuration. • Example: Hydrogen (AtomicNumber 1) is 1s1 • Helium (Atomic Number 2) is 1s2 • Neon (Atomic Number 10) is 1s2 2s2 2p6

  20. Rule 1: Aufbau Principle • AufbauPrinciple:Electrons always fill orbitals of lower energy first. 1s is filled before 2s, and 2s before 2p. • Electrons are placed in increasing energy: • 1s<2s<2p<3s<3p<4s<3d<4p<5s<4d<5p<6s<4f<5d<6p<7s<5f<6d<7p

  21. An Easier Way to Use the Aufbau Principle:

  22. Rule 2: Pauli Exclusion Principle This discovery was made in 1925 by German chemist Wolfgang Pauli → • No two electrons within a particular atom can have identical quantum numbers. If two electrons do occupy the same orbital, they must have opposite spin.

  23. Rule 3: Hund’s Rule • When an electron joins an atom and has to choose between two or more orbitals of the same energy, the electron will prefer to enter an empty orbital rather than one already occupied. • As more electrons are added to the atom, electrons will half-fill orbitals of the same energy before pairing with existing electrons to fill orbitals. http://chematomicstructure.wikispaces.com/7a.+Electronic+Configuration+-+3+Rules

  24. Practice:Writing out Electron Configuration of Silicon Step A Step B Step C Step D Step E 1s22s2p63s23p2

  25. Now, You Practice! • Write out the electron configuration of Potassium:

More Related