1. Pre-Class Activity Pass around the box. “Examine” what is inside without opening the box. Try to figure out what is in the box. What observations did you make to come to this conclusion? Section 5.1 Homework Quiz Friday

2. Unanswered Questions Rutherford’s model did not address the following questions: 1.What is the arrangement of electrons in the atom? 2.What keeps the electrons from converging on the nucleus? 3.What accounts for the differences in chemical behavior among elements?

3. Chemical Behavior Related to the Arrangement of Electrons In the early 1900’s scientist observed that when they heated elements in a flame, the elements emitted colored light. Analysis of the light led scientist to determine that the arrangement of electrons in an atom is related to the elements chemical behavior.

4. Understanding Light – Characteristics of Waves C A 1 sec D E Using the word bank below, match the words to the lettered parts of the diagram B = 3/s FrequencyAmplitudePeak TroughWavelength

5. Waves Characteristics Wavelength ( ) the shortest distance between two equivalent points on a wave. ( unit of measure: meters, centimeters, or nanometers) Frequency (v ) the number of waves that pass a given point per second. (unit of measure: Hz, waves/second, /s, s -1 ) c v Waves c= v Speed of Light ( c )

6. Relationship Between Wavelength and Frequency c= v Frequency Wavelength

7. Problems What is the wavelength of a wave that has a frequency of 7.8 x 10 10 Hz? What is the frequency of a wave that has a wavelength of 6.5 x 10 -4 cm?

8. Pre-Class Activity Which of the following waves has the longest wavelength? Which has the greatest frequency? 5.1 Quiz tomorrow, Atomic Emission Spectroscopy Lab on Thursday

9. Electromagnetic Spectrum Electromagnetic radiation – Form of energy that exhibits wave-like behavior. Electromagnetic spectrum – encompasses all forms of electromagnetic radiation. The only differences between each form is their characteristic wavelengths and frequencies. White Light – a continuous spectrum of colors, each with a unique wavelength and frequency.

10. Particle Nature of Light Concept developed to explain why only certain frequencies of light are given off by heated objects and why metals will eject electrons from their surface when certain frequencies of light is shine on them.

11. Max Planck and the Quantum By examining the light emitted by glowing objects, Planck concluded that matter can gain and lose energy in only small specific amounts called quanta Quantum – the minimum amount of energy that can be gained or lost by an atom. Temperature Energy Temperature Energy

12. Planck Continued E quantum =hv E = energy h = Planck’s constant=6.626 x 10 -34 J V = frequency Relationships between wavelength, frequency and Energy Frequency Wavelength Energy Frequency Energy Wavelength

13. Photoelectric Effect Electrons are emitted from a metal’s surface when light of a certain minimum frequency of 1.14 x 10 15 Hz shines on the surface. (Conversion of light energy to electrical energy)

14. Einstein’s Contribution E photon =hv Explained that electromagnetic radiation behaves both like waves and particles. Photon – a particle of electromagnetic radiation with no mass that carries a quantum of energy.

15. Bohr Model of the Atom Lowest allowable energy level = ground state Electrons move around the nucleus in only certain allowable circular orbit Each orbit is associated with a certain amount of energy The larger the orbit (farther away from the nucleus) the greater the energy

16. Line Spectra

17. Electrons as Waves (Louis de Broglie) Each energy level represents multiples of a whole wavelength de Broglie proposed the idea that all particles have wavelengths

18. Heisenberg Uncertainty Principle Premise: It is impossible to make any measurement on an object without disturbing the object. The position of electron can be determined by shooting photons at electrons in an atom It is impossible to know precisely both the velocity and position of a particle at the same time

19. Quantum Mechanical Model Like Bohr, electrons energies are limited to a certain value Unlike Bohr, electrons do not travel in prescribed paths. Electrons travel within a particular volume of space surrounding the nucleus

20. Organization of Electrons in an Atom Energy Levels n=1,2,3,4,5,6,7  Energy levels increase as distance from the nucleus increases Sublevels (s,p,d,f)  Corresponds to the shape of the atoms orbitals

21. Ground-State Electron Configuration Arrangement of electron in an atom at rest  The lowest energy arrangement is the most stable Three Rules that govern how electrons can be arranged in an atom:  Aufbau Principle (each electron occupies the lowest energy orbital available)  Pauli Exclusion Principle (A maximum of two electrons can occupy a single orbital)  Hund’s Rule (Electrons with the same spin must occupy each orbital before electrons with opposite spins can occupy each orbital)

22. Orbital Filling Diagram Electron Configuration Fluorine (atomic #9) 225 Orbital Diagram

23. Nobel Gas Configuration To shorten the amount of writing required to represent the configuration of elements with large numbers of electrons, the symbol for the noble gas that directly precedes an element can represent all of the electrons up to that point. 1s 2 2s 2 2p 6 3s 2 3p 6 4s 2 [Ar] 4s2 Long HandShort Hand

24. Lewis Dot Structures Lewis dot structures represent an atoms valence (outermost) electrons that are involved in chemical bonding Writing Lewis Dot Structures 1.Determine the electron configuration of the atom 2.Identify all those electrons that exist at the highest numbered energy level 3.Write the symbol for the atom and surround it with the number of electrons identified in #2, making sure to place one electron in each of 4 locations (top, bottom, left, right) before adding additional electrons up to a total of 8 electrons. Ca: 1s 2 2s 2 2p 6 3s 2 3p 6 4s 2 Ca Lewis Dot Structure