Chapter 4 Spectroscopy & Arrangement of Electrons(e - )

Properties of Light Electromagnetic Radiation- form of energy w/ wavelike properties as it travels through space Electromagnetic Spectrum- classifies as electromagnetic radiation based on wavelength( ) and frequency(v)

The Electromagnetic Spectrum Gamma Rays X-Rays UV-Rays Infrared Microwave Radio Long-Wave Hz Hz Hz Hz 10 9 Hz Hz 10 5 Hz Short Wavelength Long Wavelength High Frequency Low Frequency 400nm 450nm 500nm 550nm 600nm 700nm Violet Blue Green Yellow Orange Red Visible Light v

Frequency(v)- # of waves that pass a given point in a specific time Wavelength( )- the distance between corresponding points on a wave (ex: peak to peak)

Speed of Light(c) Speed of Light = wavelength x frequency c = v Speed of Light is a Constant: c = 3.0 x 10 8 m/s Frequency and Wavelength are Inversely Proportional : v

Photoelectric Effect When light particles(photons) collide with a metal, the photons knock electrons (e - ) loose These electrons move toward the positive terminal creating an electric current(electricity)

Light as Particles Light has wave and particle properties (Max Planck) A Quantum of Energy- the minimum quantity if energy that can be lost or gained by an atom E = hv E = energy of a photon (J) h = Planck’s constant (6.626 x JS) v = frequency (Hz) Photons- particles of light carrying q quantum of energy

Dual Wave-Particle Nature of Light Light has both the properties of waves and of particles –Wave Properties: light can be bent as it passes through objects –Particle Properties: photons have mass and exert force on other objects

Emission Spectra Emission spectra are fingerprints of an atom Every atom gives off different colors from the visible light spectrum when they release absorbed energy

Energy States of Atoms/Electrons Ground State- lowest energy level of an electron within an atom Excited State- a higher energy level within an atom that an electron may exist in –Energy must be absorbed for an electron to go from ground to excited state –Energy is given off as visible light when an atom returns to ground state –Every atom gives off a unique spectrum based on the movement of it’s electrons

Energy of a Photon E photon = E final - E initial E2E2 E1E1 E 2 – E 1 = E photon = hv Excited State Ground State

Bohr Model of the Hydrogen Atom Bohr’s model indicates that as atoms absorb energy their electrons move to higher energy levels When the absorbed energy is given off as visible light the electrons return to their ground state

Bohr Model

Quantum Model of the Atom The Bohr model was more accurate than previous models but was only completely accurate for Hydrogen, other elements did not behave exactly as Bohr predicted The Quantum model was later developed based on work of many scientists including Schrodinger, Heisenberg, & Einstein

Quantum Model of the Atom

Electrons as Waves Louis deBroglie proved that electrons had wave properties by showing that electrons could produce interference patterns like sound and light waves Passing electrons through a crystal also caused the stream of electrons to bend like light waves do

Interference Patterns

Heisenburg Uncertainty Principle This principle states that it is impossible to determine simultaneously both the position and velocity of an electron This theory led to the concept of the electron cloud

Quantum Theory This theory describes mathematically the wave properties of electrons based on probability –Electrons are not in set energy levels –Electrons are in 3D orbits around the nucleus called orbitals

Orbitals Orbitals are 3D regions around the nucleus of an atom that indicate the probable location of an electron

Quantum #’s Quantum #’s are used to specify the properties and location of electrons in orbitals around the nucleus There are 4 quantum #’s, each is represented by a letter : n, l, m, & s

Principle Quantum #(n) The principle quantum number indicates the main energy level of an electron and it’s distance from the nucleus n=1 : ground state, e - close to nucleus n=7 : excited state, e - further from nucleus * As the n value increases so does the energy of the e -, and the distance from the nucleus

Angular Momentum Quantum #( l ) The angular momentum quantum # indicates the shape of an orbital Each l value has a corresponding shape l = 0 : s-shape, sphere orbital around nucleus - an s-orbital can hold up to 2 electrons l = 1 : p-shape, 2 lobes on either side of nucleus - an atom can have 3 p-orbitals, one in each plane(x,y,z) - a p-orbital can hold up to 6 electrons

Angular Momentum Quantum #( l ) l = 2 : d-shape, 4 lobes “clover” around nucleus –An atom can have up to 5 d-orbitals –A d-orbital can hold up to 10 electrons

Angular Momentum Quantum #( l ) l = 3 : f-shape, 8 lobes around nucleus –An f-orbital can hold up to 16 electrons

Magnetic Quantum #(m) The magnetic quantum # indicates the orientation of an orbital around the nucleus

Spin Quantum #(s) The spin quantum number indicates the direction that an electron is spinning, either clockwise or counterclockwise s = +1/2 : clockwise spin s = -1/2 : counterclockwise spin * 2 electrons in the same orbital must have opposite spins

Quantum #’s n – principle quantum # distance from nucleus / main energy level l – angular momentum quantum # shape of orbital m – magnetic quantum # orientation of orbitals around nucleus s – spin quantum # direction of e - spin around nucleus

Pauli Exclusion Principle The Pauli Exclusion principle states that no 2 electrons in the same atom can have the same combination of 4 quantum #’s This means that no 2 electrons could be in the same place at the same time

Electron Configuration Notation 1s 2 2s 2 2p 6 3s 2 3p 6 4s 2 3d 10 4p 6 5s 2 4d 10 5p 6 6s 2 … Ex: Carbon = 6 electrons C = 1s 2 2s 2 2p 2 Ex: Sodium = 11 electrons Na = 1s 2 2s 2 2p 6 3s 1 Principle Quantum # or Energy Level Angular Momentum Quantum # or Orbital Shape # of Electrons in Orbital

Noble Gas Notation(Shortcut) To eliminate repetitive electron configuration for elements with large #’s of electrons the symbol of a Nobel Gas can be substituted for a portion of the electron configuration Ex: K = 1s 2 2s 2 2p 6 3s 2 3p 6 4s 1 [Ar]4s 1 Ex: Zn = 1s 2 2s 2 2p 6 3s 2 3p 6 4s 2 3d 10 [Ar]4s 2 3d 10

Orbital Notation Orbital notation uses boxes and arrows to indicate electrons in orbital by energy level Aufbau Principle- an electron will occupy the lowest possible energy level that can hold it Hund’s Rule- orbitals of equal energy will each receive one electron before they receive a second electron

Order of Atomic Sublevels Orbital Notation: P: O: * Each p-orbital gets 1 e - before it gets a 2 nd e -

Examples N : Si : Fe : Mg :

e - Configurations 1s 2s 3s 4s 5s 6s 7s7p 3p 4p 5p 6p 2p 3d 4d 5d 6d 4f 5f

Flame Test Lab CompoundColor Description NaCl NaNO 3 Sr(NO 3 ) 2 Ca(NO 3 ) 2 Ba(NO 3 ) 2 KNO 3 Cu(NO 3 ) 2 CuSO 4 LiNO 3 UnknownColor Description Compound A B C D E F

Calculations CompoundFlame Color λ (nm) ν (Hz) E (J) LiNO 3 Ba(NO 3 ) 2 Ca(NO 3 ) 2 General Equations: Speed of Light: c = λ ν Energy: E = h ν c = 3.0 x 10 8 m/sh = x Js

END OF CHAPTER 4 NOTES !!!