2.  A. General Information › 1. electron behavior has been studied through light › 2. remember, light IS radiant energy › 3. originally considered to be wave energy alone › 4. in the 1900’s, scientists determined light behaved like a particle  So it is both!

3.  B. Waves 1. Light waves are electromagnetic waves Called electromagnetic radiation (ER) X-rays, gamma rays, and radio waves are, also, a part of ER 2. Electromagnetic waves consist of electric and magnetic fields oscillating at right angles. a)All waves are described by 4 characteristics

5.  B. Waves 2. Electromagnetic waves consist (CONT.) a. all waves are described by 4 characteristics 1.Amplitude – the height of the wave measured from its origin to its peak or crest The brightness or intensity of light is dependent on this part of the wave. 2.Wavelength – the distance between successive crests The distance traveled in a full cycle Visual light has a range between 400 to 750 nm (10 -9 )

6.  B. Waves 2. Electromagnetic waves consist (CONT.) a. all waves are described by 4 characteristics 3. Frequency – how fast the wave oscillates up and down (during a given time, usually 1 second)  The unit is cycle/s or hertz (Hz)  1 Hz = 1 cycle/s  Ex. FM Radio 93.1 MHz = 93.1 x 10 6 cycle/s  Ex. Visual light is between 4 x 10 4 cycle/s and 7 x 10 14 cycle/s

8.  B. Waves 2. Electromagnetic waves consist (CONT.) a. all waves are described by 4 characteristics 4. Speed – which is a constant value Called “the speed of light” = 3/00 x 10 8 M/S. This creates a direct relationship between wavelength and frequency. The shorter the distance, the greater the oscillations. The longer the distance, the fewer the osscilations.

9.  B. Waves 2. Electromagnetic waves consist (CONT.) b. The relationship between frequency and wavelength is a mathematical expression.  λ (lambda) = wavelength, V (nu) = frequency, c = speed of light  λ = v * c  Ex. Helium – neon laser has a wavelength of 633 nm…v = ?  4.74 x 10 -14 s -1

10.  C. Electromagnetic Spectrum 1. Prisms separate light into the different wavelengths a.A rainbow is all of the light in the visible spectrum (ROY G BIV) Violet has the shortest wavelength, Red has the longest wavelength b.Visible light constitutes of a very small portion of the electromagnetic spectrum. c.The rest of the electromagnetic spectrum is invisible to the eye. Consists of Gamma rays, X-rays, UV, visible, infra red (IR), Microwaves, TV waves, and Radio From smallest wave to longest wave

12.  A. General Information › 1. Hot objects emit electromagnetic waves (why?)  a. first emits heat (IR energy/light)  b. begins to glow (Red to yellow to white for metal)  Electric stove tops › 2. Barium and Strontium emit green and red colors (why?) › 3. Gases give off specific colors of light when heated

13.  B. Planck’s Theory 1. Max Planck theorized the spectrum of radiation emitted changes with temperature 2. Theorized energy emitted or absorbed is restricted to “pieces” of particular size” 3. Proposed – There is a fundamental restriction on the amounts of energy that an object emits or absorbs, which are called quantum ( meaning fixed amount)  Derived from the concept of the relationship between frequency (v) and the energy (E) with which it is associated  Plank determined the energy constant, known as Plank’s constant (h) with a value of 6.6262 x 10 -34 J/s.

14.  B. Planck’s Theory 4. Plank’s equation is E = H * v 5. Quantum of energy of extremely small, so it looks like a continuous climb

15.  C. Photoelectric Effect 1. Electrons are ejected from metal when light is shined on it. 2. A minimum frequency of light is needed to release electron Ex. Sodium metal wont release electrons with red light, but will with violet light 3. Light consists of quanta of energy that behave like tiny particles of light. Called photons Photon energy is equal to Plank’s energy

18.  A. Line Spectra 1. A line spectrum is a spectrum of colors created from a prism 2. Elements emit light when they are vaporized 3. Each element has a unique line spectrum An atomic fingerprint 4. Each element when placed under a flame appears as a color Salt – yellow, because of sodium Lithium – red Potassium – blue w/ red Neon – red Nitrogen - orange

20.  B. Bohr Model (Neils Bohr) 1. Used Rutherford’s planetary model to help explain element emissions 2. Orbitals around the nucleus were based on quanta and given a quantum number Lowest level, n = 1 (Ground state) 3. When the electron absorbs enough energy, it will jump to the next energy level (excited state) n = 2, 3, 4, etc. 4. Light is emitted as the electron “falls” back to the ground state

22.  C. Matter Waves 1. Louis De Broglie theorized that matter has a dual nature. Believed matter should have wave like behavior and exhibit wavelengths Called it matter waves Came up with a mathematical formula relating the mass and velocity of a moving particle and its possible wavelength. This finding was used to create the electron microscope. For waves to be seen from objects the mass must be very small.

23.  D. Heisenberg’s Uncertainty Principle 1. Werner Heisenberg proposed that the position and the momentum of a moving object cannot be simultaneously measured and known exactly. Uncertainty Principle 2. Hard to predict where a particle will be in the future.

24.  A. General Information 1. 3 known concepts shape the new look a.Energy is quantized b.Electrons exhibit wavelike behavior c.Impossible to know exactly where an electron is in space 2. These 3 concepts lead to the quantum- mechanical model

25.  B. Probability and Orbitals 1. General Information a.Consider an electrons place around the nucleus as a blurry cloud b.The cloud’s density is greater where there is a higher probability that the electron is present Called electron density

26.  B. Probability and Orbitals 1. General Information c. An atomic orbital is a region around the nucleus of an atom where an electron with a given energy is likely to be found Orbitals have characteristic shapes (not energy) Draw orbitals based on where they are likely to be located 90% of the time Different orbitals are designated by different letters S, p, d and f S = spherical shaped P = dumbbell shaped D & f = complex

28.  B. Probability and Orbitals 2. Orbitals and Energy a.The Principle Energy Levels in an atom are designated by the quantum number (n) n is the principle quantum number

29.  B. Probability and Orbitals 2. Orbitals and Energy b. The energy of the electron increases as n increases (1, 2, 3, 4, 5, 6) Each energy level is divided into one of more sublevels The number of sublevels in each principle energy level = quantum number For example 1 = 1 sub, 2 = 2 sub, 3 = 3 sub, etc The sublevels are indicated by a letter. For example n = 1; 1s n = 2; 2s 2p n = 3; 3s 3p 3d N = 4; 4s 4p 4d 4f

30.  B. Probability and Orbitals 2. Orbitals and Energy b. The energy of the electron increases as n increases (1, 2, 3, 4, 5, 6)  The number of orbitals in each sublevel is always equals the quantum number  n=1; 1s : 1 spherical orbital  n=2; 2s : 1 spherical orbital (larger than 1s) 2p: 3 bell shaped orbitals  n=3; 3s I spherical orbital (larger) 3p: 3 bell shaped orbitals (larger than 2p) 3d: 5 complex (d) orbitals  n=4; 4s: 1 spherical orbital (larger() 4p: 3 bell shaped orbitals (larger) 4d: 5 complex orbitals (larger than 3d) 4f: 7 complex orbitals

32.  B. Probability and Orbitals 3. Electron Spin a.Electrons spin on their access (2 ways only) Can spin clockwise and counterclockwise b.Spinning charges create magnetic fields Clockwise in N ↑ Counterclockwise is N ↓ c.Can have parallel spins or opposite spins If opposite spins, cancel the magnetic pull If parallel, they create magnetic effect d.Pauli exclusion principle Each orbital in an atom can hold only 2 atoms with opposite spins.

34.  A. General Information 1. Electron configuration is the distribution of electrons among the orbitals of an atom 2. Electron configurations describe where the electrons are found and what energies they possess 3. Electron configurations of atoms are determined by distributing the atom’s electrons among levels, sublevels, and orbitals based on a set of stated principles.

35.  B. Determining Electron Configurations 1. Easy once you learn the energy levels of the orbitals within each principle energy level, the s-sublevel is the lowest level. 2. When electrons populate the lowest energy orbitals, they are in the ground state.

37.  B. Determining Electron Configurations 3. The electron locations can be predicted by using the Aufbau principle, the Pauli Exclusion principle, and Hund’s Rule a.Aufbau – electrons are added one at a time, to the lowest energy orbitals available Until all electrons are accountable b. Pauli Exclusion principle – An orbital can hold a maximum of 2 electrons To occupy the same orbital, the electrons must have opposite spins (called paired electrons) c. Hund’s – Electrons occupy equal-energy orbitals so that a maximum number of unpaired electrons result.

41.  B. Determining Electron Configurations › 4. Arrows represent electrons, boxes represent orbitals called orbital diagrams (show electrons in orbitals)  Ex. 6 C : 1s 2s 2p  ↑ = counterclockwise spin  ↓ = clockwise spin e. electron configurations are created from electron diagrams (and vice versa) ex. 6 C = 1s 2 2s 2 2p 2 Exponents give you the number of electrons in each energy level. ↑↓ ↑ ↑

45.  C. Exceptions to Aufbau’s Principle 1. Some elements don’t follow the rule a.“They are interesting” Ex. Cr-23, Cu-29 Expected: Cr-23: 1s 2 2s 2 2p 6 3s 2 3p 6 4s 2 3d 4 Actual: 1s 2 2s 2 2p 6 3s 2 3p 6 4s 1 3d 5 Expected: Cu-29: 1s 2 2s 2 2p 6 3s 2 3p 6 4s 2 3d 9 Actual: 1s 2 2s 2 2p 6 3s 2 3p 6 4s 1 3d 10 b.Cause by interactions of electrons in orbitals with very similar energies