Electron Configurations Chapter 5 Electron Configurations

Light - Light has properties of both waves and particles Properties of Waves Amplitude – height of a wave Wavelength () – distance between waves or distance between two corresponding points Frequency () – how fast the wave oscillates or number of waves per second Speed – Speed of light (c) = 3.00 x 108 m/s All waves (electromagnetic, water) are described using these four properties

Wavelength Amplitude Frequency : unit = cycles/sec = 1/s = s-1 = Hertz

Speed of Light (c) - Light moves at a constant speed (c = 3.00 x 108 m/s) - Creates a relationship between frequency and wavelength wavelength () = speed of light (c) / frequency() which can be written  = c /  Example Calculate the wavelength of visible light with a frequency of 5.5 x 1014 s-1. Given:  = 5.5 x 1014 s-1 = 3.00 x 108 (m/s) / 5.5 x 1014 (1/s) = 5.5 x 10-7 m c = 3.00 x 108 m/s Find:  Equation:  = c/ 

Electromagnetic radiation is a form of energy that exhibits wavelike behavior Electromagnetic spectrum (EM) encompasses all the forms of radiation Atomic emission spectrum of an element is the set of frequencies of the electromagnetic waves emitted by an atom

Electromagnetic Spectrum 700 nm 400 nm Figure 4-7, Page 129

Energy E= hν h= 6.626x10^34 Joules/second

Quantum Theory: Plank’s Theory - There is a fundamental restriction on the amount of energy that an object emits or absorbs - Energy is emitted or absorbed in “Pieces” of particular size called quantum: is the minimum amount of energy that can be gain or lost by an atom - Quantum theory contradicts classical theory which states, an object can absorb or emit any amount of energy, so the energy it can posses forms a continuum of values Quantum Theory answers the question: Why electromagnetic radiation is emitted by hot objects

Essence of Plank’s Theory - provides a relationship between frequency () and energy (E) E = h  where h = plank’s constant = 6.626x10-34 J.s  = frequency = Hertz (s-1) E = energy = Joule (J) - showing that, energy is quantized; restricted to certain quantities Example What if your car was quantized?

- Einstein used this idea of quantized energy to explain a - Einstein used this idea of quantized energy to explain a phenomenon called the photoelectric effect Photoelectric Effect e- are ejected from the surface of a metal when light shines on the metal Einstein concluded that light is quantas of energy called photons Photons (quantas of light energy) strike the surface of the metal Step 1 The energy is transferred to the e- in a metal atom Step 2

Step 3 The e- “swallows” the entire photon, OR The e- “swallows” none of the photon * Energy is consumed all or nothing* Step 4 (a) If there is NOT enough energy in a particular photon: e- stays put No “roll over” plan. Energy does NOT collect. No mater how many times you hit with the same energy photon (b) If there is enough energy in a particular photon: e- is ejected from surface Conclusion The frequency, and therefore energy, is important in a photon, not the intensity of the light Figure 4-12, Page 133

Vocabulary Line Spectrum – A spectrum that contains only certain colors, or wavelengths Ex. Street lights, neon lights Continuous – Spectrum A spectrum that contains all colors, or wavelengths, which fade into eachother Ex. Electromagnetic spectrum, sunlight, filament bulb Quantum number (n) - Indicates an energy level Ground State (n=1) - Lowest energy level; orbital closest to the nucleus Excited State (n=2, ...) - energy levels above ground state achieved by electrons thru absorbing energy Atomic orbital – the region around the nucleus where an electron with a given energy is likely to be found Heisenberg’s Uncertainty -Principal The position and momentum of an e- cannot be simultaneously be known

ENERGY E = hυ h=6.636x10^-34 Max Planck

Bohr’s Model of the atom Quantum # (n)- Indicated energy level or orbital Neils Bohr n = 1 + n = 2 The probability of finding e- in certain regions of an atom is described by an orbital n = 3 n = 4 Orbitals have characteristic shapes, size and energies, but do NOT tell how e- move The larger the orbital, the further the e- are from the nucleus

- There are four types of orbitals: s, d, p, f

s p d f Electron Configurations 1 2 1 2 3 4 5 6 7 1 2 3 4 5 6 1 2 valence e- = outermost e- 1 2 3 4 5 6 7 s 1 2 3 4 5 6 Columns (top to bottom) = group/family Rows (left to right) = period 2 p 3 1 2 3 4 5 6 7 8 9 10 3 d 4 4 5 5 6 6 7 1 2 3 4 5 6 7 8 9 10 11 12 13 14 4 f energy levels 5

Quantum model

Electron Spin - e- spin either clockwise or counterclockwise which creates a magnetic field each orbital in an atom can hold up to 2e- that must have opposite spin ( and ) Pauli Exclusion Principle – - an orbital with 2e- of opposite spin are said to have a pair of e- Sublevels # Orbitals Max # of e- s 1 2 p 3 6 d 5 10 f 7 14 Figure 4-24

Electron Configurations Definition Distribution of e- among orbitals; describes where e- are and what energy they have Aufbau Principle – e- are added one at a time to the lowest energy level available Hund Rule – e- occupy equal energy orbitals so the maximum number of unpaired e- result 3s _____ 2p _____ _____ _____ 2s _____ 1s _____          Figure 4-24

Electron Configurations Definition Distribution of e- among orbitals; describes where e- are and what energy they have Aufbau Principle – e- are added one at a time to the lowest energy level available Hund Rule – e- occupy equal energy orbitals so the maximum number of unpaired e- result 3s _____ 2p _____ _____ _____ 2s _____ 1s _____          Figure 4-24

Electron Configurations Definition Distribution of e- among orbitals; describes where e- are and what energy they have Aufbau Principle – e- are added one at a time to the lowest energy level available Hund Rule – e- occupy equal energy orbitals so the maximum number of unpaired e- result _____ _____ _____ _____ _____ 1s 2s 2px 2py 2pz 1s2 2s2 2p6         

1s2 2s2 2p6 Orbital Last level of energy Number of electrons in a particular orbital Level of energy This particular atoms has 2 levels of energy 10 electron total 2 + 2 +6 8 electrons in the last level of energy 2 +6

How To Do Energy Diagrams Step 1 Determine atomic number Step 2 Determine number of e- - if neutral: # e- = atomic # - if ion: # e- = atomic # - charge Step 3 Fill in atomic orbital energy diagram using following principles - Pauli: 2e- max per orbital with opposite spin - Aufba: 1e- at a time filling lowest energy levels first - Hund: max # of unpaired e- Step 4 Determine number of paired and unpaired e- Step 5 Write e- configuration

Example. Determine the e- configuration of Ne by using an Example Determine the e- configuration of Ne by using an atomic orbital energy diagram 3s _____ 2p _____ _____ _____ 2s _____ 1s _____ Atomic # = 10 # e- = 10 – 0 = 10 Paired = 5 Unpaired = 0         e- config: 1s22s22p6   Example Determine the e- configuration of S by using an atomic orbital energy diagram 4s _____ 3p _____ _____ _____ 3s _____ 2p _____ _____ _____ 2s _____ 1s _____ Atomic # = 16 # e- = 16 – 0 = 16 Paired = 7 Unpaired = 2       e- config: 1s22s22p63s23p4          

Example. Determine the e- configuration of Fe+ by using an Example Determine the e- configuration of Fe+ by using an atomic orbital energy diagram 3d _____ _____ _____ _____ _____ 4s _____ 3p _____ _____ _____ 3s _____ 2p _____ _____ _____ 2s _____ 1s _____        Atomic # = 26 # e- = 26 – 1 = 25 Paired = 10 Unpaired = 5                 e- config: 1s22s22p63s23p64s23d5   Example Determine the e- configuration of F- by using an atomic orbital energy diagram 3s _____ 2p _____ _____ _____ 2s _____ 1s _____ Atomic # = 9 # e- = 9 – -1 = 10 Paired = 5 Unpaired = 0         e- config: 1s22s22p6  

Order of Energy Levels from Lowest (1s) to Highest Group 1 – Alkali Metals Group 2 – Alkaline Earth Metals Group 1- 10 – Transition Metals Group 3 – Boron Group Group 4 – Carbon Group Group 5 – Nitrogen Group Group 6 – Oxygen Group Group 7 – Halogens Group 8 – Nobel Gas Period 4f– Lanthanides Period 5f– Actinides Start with 1s and read up the arrow