1. Electrons in Atoms By: Ms. Buroker

2. Okay … We now know that an element’s identity lies in its number of protons … but there is another particle which is very important as well … The electron Electrons control behavior and reactivity in substances.

3. The Electromagnetic Spectrum Visible light is a kind of electromagnetic radiation … a form of energy that exhibits wavelike behavior as it travels through space.

4. Electromagnetic Radiation We measure electromagnetic radiation in two ways … 1.) Wavelength 2.) Frequency When multiplied together … these two thing ALWAYS equal the speed of light!!! C= 3.00 x 10 8 m/s

5. Wavelength ( ) Wavelength is the distance from one wave crest or wave trough to the next

6. Electromagnetic Spectrum Wavelength is measured in meters … but more commonly, nanometers.

7. Frequency (  ) Frequency is defined as the number of waves that pass a given point in a specific time … usually, a second. Typically expressed as a Hertz (Hz) … one wave per second

8. Remember What We Said … All Wavelengths and their frequency’s when multiplied are equal to the speed of light; so … C = 

9. Let’s Re-Visit This … Violet … shortest wavelength and highest frequency Red … longest wavelength and lowest frequency

10. The Particle Nature of Light Matter can loose energy only in small, specific amounts called quanta … so a quantum is the minimum amount of energy that can be gained or lost by an atom. Energy Plank’s Constant 6.626 X 10 -34 J.s Frequency E quantum = h

11. The Photoelectric Effect Einstein suggested that electromagnetic radiation can be viewed as a stream of particles called photons. E photon = h = hc Photoelectric Effect: refers to the phenomenon in which electrons are emitted from the surface a metal when light strikes it.

12. Photoelectric Effect Continued 1.) Each metal has a threshold frequency. 2.) For light with frequency lower than the threshold, no e- are emitted. 3.) For light with frequency greater than threshold, the number of e- emitted increases with the intensity of the light. 4.) For light with frequency greater than the threshold frequency, the kinetic energy, of the emitted e- increases linearly with the frequency of light.

13. Let’s Practice!! Tiny water drops in the air disperse the white light of the sun into a rainbow. What is the energy of a photon from the violet portion of the rainbow if it has a frequency of 7.23 x 10 14 Hz? E photon = h E photon = (6.626 x 10 -34 J.s)(7.23 x 10 14 /s) E photon = 4.79 x 10 -19 J

14. Atomic Emission Spectra The set of frequencies of the electromagnetic waves emitted by atoms of an element. * They are unique to the element!

16. Quantum Theory and the Atom Bohr Model of the Atom * Working with the hydrogen atom, he proposed that the hydrogen atom has only certain allowable energy states. Lowest allowed energy state = ground state Atoms gain energy = excited state

17. Bohr Model Continued … Electrons moving around the nucleus in only certain allowed circular orbits

18. Bohr Model Continued … Assigned each orbital a quantum number, n Smaller the orbital = lower energy Higher orbitals = higher energy

19. de Broglie Equation de Broglie predicted that all moving particles have wave characteristics. = h mv Mass Velocity

20. Heisenberg Uncertainty Principle States that it is impossible to know the precisely both the velocity and position of a particle at the same time.

21. Hydrogen’s Atomic Orbitals n = Principle Quantum Number Energy sub levels = s, p, d, f Shape = s: spherical p: dumbbell shaped d,f: don’t all have the same shape

22. S orbitals 2 e - p orbitals 6 e - d orbitals 10 e - f orbitals 14 e -

23. In Summary Principle Quantum Number (n) Sublevels (types of orbitals) Present Number of Orbitals related To sublevel 1s1 2spsp 1313 3spdspd 135135 4spdfspdf 13571357

24. Electron Configuration The arrangement of electrons in an atom. There are three main rules that govern how electrons can be arranged in an atom.

25. The aufbau principle e - occupy the lowest energy level FIRST!

26. The Pauli exclusion principle A maximum of two e - may occupy a single atomic orbital, but only if the electrons have opposite spins! orbitalTotal # of e- s2 p6 d10 f14

27. Hund’s rule Single e - with the same spin must occupy each equal energy orbital before additional e - with opposite spin can occupy the same orbital.

28. Atomic Orbitals and Quantum Numbers n= The principle quantum # (1, 2, 3, ….) * relates to the size and energy of the orbital l = The angular momentum quantum number (0 to n-1) l = 0  s orbital l = 1  p orbital l = 2  d orbital l = 3  f orbital

29. Atomic Orbitals and Quantum Numbers M l = magnetic quantum number (l to –l) * related to the orientation of the orbital in space relative to the other orbitals in the atom. M s = Electron spin quantum number (+1/2 or -1/2)

30. Valance Electrons Valence electrons are those electrons that are in the outer most energy level of an atom … it is these electrons that are responsible for the reactivity of an element. Example: Write the electron configuration for Carbon. 1s2 2s2 2p2 … the outer most energy level for Carbon is n=2, right? So to find out how many valance electrons carbon has, you simply count how many electrons are in level 2. Carbon has 4 valance electrons!

31. Believe it or not … you can look at the periodic table and determine how many valence electrons an atom has and how that affects the properties of that element.

32. All elements with the exception of the transition metals are called the representative elements. The representative elements follow the following rules … The period number tells you the energy level the valence electrons are in!! The group number tells you how many valence electrons there are!!