Electrons

1. Electrons

2. Wave model – scientist say that light travels in the form of a wave

3. Electromagnetic Radiation – is a form of energy that exhibits wavelike behavior as it travels through space

5. All electromagnetic waves travel at a speed of 3.00 x 108 m/s in a vacuum • 3.00 x 108 m/s = the speed of light

6. Electromagnetic spectrum encompasses all forms of electromagnetic radiation showing their different wavelengths and frequencies

8. Because all electromagnetic waves travel at the same speed you can use a formula to calculate the wavelength or frequency • c = ƒλ c = speed of light ƒ = frequency λ = wavelength

9. Planck – says light travels as a particle

10. Quantum Concept • Quantum is the minimum amount of energy that can be gained or lost by an atom • This describes that matter can gain or lose energy only in small, specific amounts

11. Planck showed that energy of quantum is related to frequency of the emitted radiation mathematically: Equantum= hv E = energy h = Planck's constant (6.626 x 10-34J) v = frequency

12. According to this theory, for a given frequency, matter can emit or absorb energy only in whole-number multiples of hv

13. Photoelectric Effect – • Photoelectrons are emitted from a metal’s surface when light of a certain frequency shines on the surface • Calculators

14. Wave model says that given enough time all light with different types of frequencies and wavelengths will accumulate and supply enough energy to eject photoelectrons from a metal • Metals will NOT eject photoelectrons below a specific frequency

15. Albert Einstein proposed that electromagnetic radiation has both wavelike and particle like natures. • A photon is a particle of electromagnetic radiation with no mass that carries a quantum of energy

16. Due to this Einstein said: Ephoton = hv E = energy h = planck’s constant V = frequency

17. Bohr’s Atomic Model: Hydrogen • Bohr says that an atom has only certain allowable energy states • The lowest energy state is called its ground state

18. When an electron gains energy it is in an excited state • Atoms are capable of having many different excited states

19. Bohr’s model showed that electrons traveled around the nucleus in a specific orbital • The smaller the electron’s orbital the lower the atom’s energy state/level • So the larger the orbital the higher the atom’s energy level

20. A quantum number was assigned to each orbital • The first orbital, the closest to the nucleus, is n = 1

21. When electron’s are at ground state, they do not radiate any energy • When energy is added from an outside source the electrons move to higher energy levels. The atom is then in an excited state

22. Electrons that are in an excited state can drop from the higher energy level to a lower energy level. As a result the atom emits a photon that is equal to the difference between the two energy levels

23. Ehigher-energy orbit – Elower-energy orbit = Ephoton • Ephoton = hv

24. Electrons can only move from one allowable orbit to another, similar to climbing a ladder one step at a time • This controls how much energy can be emitted or absorbed

25. The Quantum Mechanical Model of the Atom • Light is a wave with particle like characteristics • Electrons are particles with wave like characteristics

26. The Heisenberg uncertainty principle states that it is fundamentally impossible to know precisely both the velocity and position of a particle at the same time.

27. Heisenberg tried to locate the position and velocity of an electron by using a photon of electromagnetic wave, but the photon disturbed the electron and its position or velocity was changed.

28. The quantum mechanical model of the atom shows electrons treated as waves • It shows limits to an electron’s energy absorbed or emitted • But it does not describe the electron’s path around the nucleus

29. The quantum mechanical model assigns principal quantum numbers that indicate the relative sizes and energies of atomic orbitals not location from the nucleus.

30. As n (principal quantum number) increases the orbital becomes larger, electron is farther from the nucleus and the energy level increases

31. Principal energy levels contain energy sublevels. • Principal energy level 1 consists of a single sublevel • Principal energy level 2 consists of two sublevels • And so on

32. Sublevels are labeled s, p, d, or f according to the shapes of the atom’s orbitals. • s orbitals are spherical • p orbitals are dumbbell shaped • d and f orbitals don’t all have the same shape

34. Principal energy level 1 has 1 sublevel • 1s • The s orbitals are spherical

35. Principal energy level 2 has 2 sublevels • 2s, 2p • p orbitals consist of 3 dumbbell shaped p orbitals of equal energy • Labeled 2px, 2py and 2pz

36. Principal energy level 3 consists of 3 sublevels • 3s, 3p, 3d • Sublevel d consists of five orbitals of equal energy they have an identical shape but different orientation

38. The maximum number of orbitals related to each principal energy level is n2 • Example: • Principal energy level 1 has 1 orbital • Principal energy level 2 has 4 orbitals • Principal energy level 3 has 9 orbitals

39. The maximum amount of electrons related to each principal energy level is 2n2 • Because 2 electrons can go in each orbital • In Principal energy level 1 there can be 2 electrons • In principal energy level 2 there can be 8 electrons

40. Electron Configuration • The arrangement of electrons in an atom is called the atom’s electron configuration • Electrons in an atom tend to assume the arrangement that gives the atom the lowest possible energy • the atoms ground-state electron configuration

41. There are 3 rules, or principles that define how electrons can be arranged in an atom’s orbitals • aufbau principle • Puali exclusion principle • Hund’s rule