Electrons in Atoms

1. Electrons in Atoms Chapter 5

2. In this chapter… • Scientists pursue an understanding of how electrons are arranged within atoms • Electron arrangement plays a role in chemical behavior

3. Chemists found Rutherford’s nuclear model to be lacking because it did not begin to account for the differences in chemical behavior among various elements • Early 1900’s- scientists observed that certain elements emit visible light when heated in a flame chemical behavior

4. Wave Nature of Light • Electromagnetic Radiation- form of energy that exhibits wavelike behavior as it travels through space

5. Vocabulary to know.. • Wavelength- shortest distance between equivalent points on a continuous wave • Symbol- λ (lambda) • Unit- meters, centimeters, or nanometers (1 nm= 1x10-9m) • Frequency- the number of waves that pass a give point per second • Symbol- ν (nu) • Unit- Hertz (SI Unit)= (1/s)= (s-1) • Amplitude- the wave’s height from the origin to a crest, or from the origin to a trough

6. How are they related? • ALL electromagnetic waves, including visible light, travel at a speed of c= 3.00x108 m/s (MEMORIZE) • Speed of light= wavelength x frequency C= λν

7. Electromagnetic Spectrum • Aka EM Spectrum • Encompasses all forms of electromagnetic radiation • The only differences in the types of radiation being their wavelengths and frequencies

8. ROYGBIV

10. Calculations  • Microwaves are used to transmit information. What is the wavelength of a microwave having a frequency of 3.44 x109 Hz? • Know: C= λν • C= 3.00x108 m/s • ν = 3.44 x109 Hz • λ = ???

11. λ = c/ ν λ= 8.72 x10-2 m DON’T FORGET YOUR SIG FIG RULES!!!

12. Particle Nature of Light • Quantum Concept • Explained why colors of heated matter correspond to different frequencies and wavelengths • Max Plank- “matter can gain or lose only in small, specific amounts called quanta” • Quantum- the minimum amount of energy that can be gained or lost by an atom

13. Energy of a quantum is related to the frequency of the emitted radiation by the equation Equantum= hv • E= energy • h = Plank’s Constant (6.626x10-34J) • v= frequency • Joule (J)= SI unit for energy

14. Photoelectric Effect • Electrons, called photoelectrons, are emitted from a metal’s surface when light of a certain frequency shines on the surface • Photon- a particle of EM radiation with no mass that carries a quantum of energy Ephoton= hv

15. Atomic Emission Spectra • Set of frequencies of the electromagnetic waves emitted by atoms of the element • Example- The light of neon sign is produced by passing electricity through a tube filled with neon gas. Neon atoms release energy by emitting light.

16. An atomic emission spectrum is characteristic of the element being examined and can be used to identify that element

17. Section 5.2Bohr Model of the Atom • Proposed that the hydrogen atom has only certain allowable energy states • Ground State- lowest allowable energy state of an atom

18. Bohr’s model worked well to explain Hydrogen- however it did not explain other elements • Substantial evidence indicates that electrons do not move around the nucleus in circular orbits

19. The de Broglie equation predicts that all moving particles have wave like characteristics λ= h/mv • The Heisenburg uncertainty principal - states that it is fundamentally impossible to know precisely both the velocity and position of a particle at the same time

20. The Quantum Mechanical Model (QMM) • 1926- Austrian physicist Erwin Schrodinger used the results of Rutherford and Bohr to devise and solve a mathematical equation describing the behavior of the electron in a hydrogen atom

21. Unlike the Bohr model, the quantum mechanical model does not involve an exact path the electron takes around the nucleus • The quantum mechanical model determines the allowed energies an electron can have and how likely (probability) it is to find the electron in various locations around the nucleus • The cloud is more dense where the probability of finding an electron is high

22. Atomic Orbital- a 3D region around the nucleus describing the electron’s probable location

23. Atomic Orbitals • Principal Quantum Number (n)- indicates the relative sizes of energies of atomic orbitals • Principal Energy Levels- the major energy levels of an atom • Energy Sublevels- the energy levels contained with the principal energy level

26. Hydrogen’s First 4 Principal Energy Levels

27. Electron Arrangement in Atoms • Electrons and the nucleus interact to make the most stable arrangement possible. • Electron Configurations- the ways in which electrons are arranged in various orbitals around the nuclei of atoms

28. 3 Rules for electron configurations of atoms 1. Aufbau Principle: Electrons occupy the orbitals of lowest energy first

29. 2. Pauli Exclusion Principle: an electron orbital may describe at most two electrons • To occupy the same orbital, two electrons must have opposite spins (↓or ↑)

30. 3. Hund’s Rule- electrons occupy orbitals of the same energy in a way that makes the number of electrons with the same spin direction as large as possible.

32. Noble Gas Configuration • Method of representing electron configurations of noble gases using bracketed symbols. • Neon= [Ne] • Also used to shorten electron configurations • Sodium: #11- instead of 1s22s22p63s1 can be shortened to [Ne] 3s1

33. Exceptions to predicted configurations • Chromium- [Ar] 4s13d5 • Copper - [Ar] 4s13d10 • Illustrates the increased stability of half-filled and filled sets of s and d orbital's

34. Valence Electrons (V.E.) • Electrons in the atom’s outermost orbital's • Determine the chemical properties of an element • V.E. are used in forming chemical bonds

36. Electron Dot Structures • Consists of the element’s symbol and inner-level electrons surrounded by dots representing the atom’s valence electrons • V.E. are placed one at a time on the four sides of the symbol and then paired up until all are used

37. Examples

38.  Homework  • Page 146 • 33-35, 38-39, 43-46, 49, 52, 57-61, 64-69, 72-73, 78-82, 86-89