LIGHT and ELECTRONS

LIGHT and ELECTRONS Unit 6 Chemistry Langley

LIGHT and its PROPERTIES • Pre-1900 • Issac Newton explained light and its behavior by assuming light moved in waves • 1900 and beyond • Experimental evidence began to convince scientists that that light consists of particles (after the 1902 experiment of Max Planck) • 1905-Einstein • Dual Wave Particle Theory

LIGHT and its PROPERTIES • Wavelength: distance between two points on two adjacent waves, symbol is l (Greek symbol for lamda) • Frequency: number of waves that pass a given point in a given amount of time, symbol is n (Greek symbol nu). Units for frequency are cycles per second which SI speaking is a Hertz, Hz (Hz is also a reciprocal seconds-1).

LIGHT and its PROPERTIES • The frequency and wavelength of light are inversely proportional to each other. • As the wavelength of light increases, the frequency decreases • As the wavelength of light decreases, the frequency increases • Amplitude: Wave’s height from zero to crest or wave’s height from zero to trough (can be positive or negative) • A complete wave cycle starts at zero goes through its highest point, back through zero, reaches its lowest point, and back to zero again. • One wave cycle starts at zero and has one crest and one trough

LIGHT and its PROPERTIES • According to the Wave Model, light consists of electromagnetic waves • Electromagnetic radiation: light moving in waves through space • Radio waves, microwaves, infrared waves, visible light, ultraviolet waves, X-rays, and gamma raysElectromagnetic spectrum • Speed of light: depending on the wavlength and frequency, speed of light changes • C = ln • Speed of light in a vacuum = 3.0 x 108 m/s

SPEED of LIGHT PROBLEMS • EXAMPLE 1: • Determine the speed of light if the wavelength is 3.5 x 10-9 m/s and the frequency is 3.5 Hz.

SPEED of LIGHT PROBLEMS • EXAMPLE 2: • If light has a speed of 5.6 x 103 m/s and a frequency of 2.3 Hz, what is the wavelength.

SPEED of LIGHT PROBLEMS • EXAMPLE 3: • What is the wavelength of radiation with a frequency of 1.5 x 1013 Hz? Does this radiation have a longer or shorter wavelength than red light?

SPEED of LIGHT PROBLEMS • EXAMPLE 4: • What frequency is radiation with a wavelength of 5.00 x 10-8 m? In what region of the electromagnetic spectrum is this radiation?

PHOTOELECTRIC EFFECT(supporting work for Atomic Spectra) • The photoelectric effect is a quantum electronic phenomenon in which electrons are emitted from matter after the absorption of energy from electromagnetic radiation such as x-rays or visible light. The emitted electrons can be referred to as photoelectrons in this context. {Wikipedia.org}

PHOTOELECTRIC EFFECT (supporting work for Atomic Spectra) • Expected: Since all light is energy moving in waves, all colors of light should knock electrons off a metal • Shine different color lights on a metal • Measure the number of electrons knocked off the metal • Found that no electrons were knocked off when light was below a certain frequency

MAX PLANCK(his work used in Atomic Spectra) • German Physicists, founder of quantum theory • Studied the way light came off hot objects (diffusion of hydrogen through heated platinum) • Concluded that light comes off in small burst of particles, NOT WAVES • Quantum-minimum amount of energy that can be lost or gained by an atom • To calculate quantum/energy: E = hn • E = energy of the photon • h = Planck’s constant • n = frequency of the incident photon

ATOMIC SPECTRA • As atoms absorb energy, electrons move into higher energy levels. When the atoms release energy (lose the energy), the electron return to the lower energy levels. • The frequencies of light emitted by an element separate to give the atomic emission spectrum of the element • No two elements have the same emission spectrum

ATOMIC SPECTRA • Atomic line spectra and its existence was known before Bohr’s atomic model of hydrogen was produced. What Bohr did was explain why hydrogen had the specific frequencies it had, why it “produced/broke down” into the colors it did; it predicted the values that agreed with the experiements.

ATOMIC SPECTRA • Hydrogen Atom Line Emission Spectrum EXPECTED: Continuous spectrum of light to be given off. (Since e- are moving around nucleus randomly and using different levels of energy.) ACTUAL: Current passed through tube with Hydrogen gas. Pink light is given off. Light passed through spectrum. Found only specific frequencies of light given off.

ATOMIC SPECTRA • Lowest possible energy of the electron is referred to as its ground state • Normal location of an electron • Electrons circle the nucleus in specific orbits • If an electron absorbs energy, moves up an energy level (absorption) • If an electron gives off energy, moves down an energy level (emission)

QUANTUM MECHANICS • EINSTEIN, AGAIN!!!!!!!!!!!!!!!! • Debate between whether light is waves or particles • Einstein creates dual waves particle theory (1905) • Light is small particles (photons) that move in wave shapes • Thought electrons moved around the nucleus in wave shapes (since electrson are small particles like photons)

QUANTUM MECHANICS • Louis de Broglie: Given that light behaves as waves and particles, can particles of matter behave as waves? • Referred to the wavelike behavior of particles as matter waves • Came up with an equation that predicts all moving objects have wavelike behavior: • mv/ = h • Thanks to experiments conducted by 2 scientists, his theory was proven correctNobel Prize • Waves Waves have specific frequencies and electrons have specific orbits/energy levels • Waves and electrons can both be bent (diffraction) • Waves and electrons can both overlap and interfere with each other (interference) • Creator of Wave Mechanics

QUANTUM MECHANICS • DeBroglie’s equation combines Einstein and Planck’s equations • mv/ = h • (Anything with mass and velocity has a wavelength, so electrons have wavelengths) • DeBroglie Problems: • What is the wavelength of an electron that has a mass of 1.5 X 10-30 kg and a velocity of 2.5 X 104 m/s?

QUANTUM MECHANICS • DeBroglie Problems: • What is the velocity of an electron with a mass of 8.3 X 10-29 kg and a wavelength of 400 nm? (Hint: convert nm to m) • What is the mass of an electron with a velocity of 4.6 X 103 m/s and a wavelength of 5.6 X 10-2 meters? • What is the wavelength of an electron that has a mass of 2.8 X 10-31 kg and a velocity of 3.0 X 108 m/s?

QUANTUM MECHANICS • Heisenberg • 2 Goals in Life: • find the location of an electron • find the velocity of an electron • Problem: Electrons cannot be seen under a microscope • Only way to find an electron is to shoot a photon (particle of light) at the electron • Problem: when the photon hits the electron, it knocks the electron off course • So with this photon method, you can only know the position of an electron for a split second, but you still don’t know the velocity

QUANTUM MECHANICS • Heisenberg • DeBroglie: Tries to help Heisenberg and offers his equation  = (mv)/h • If you know mass and wavelength of an electron, equation could help you find velocity • Problem: Equation does not show location! • Equation method will only tell you velocity NOT location

QUANTUM MECHANICS • Heisenberg • Heisenberg Uncertainty Principle: It is impossible to know both the position and velocity of an electron at the same time.

QUANTUM MECHANICS • Schrodinger • Working with Hydrogen atom that only has 1 electron • Wants to find general location/area of the one electron in Hydrogen • Creates quantum theory • Quantum theory – uses math to describe the wave properties of an electron (frequency, wavelength, etc) • Once he plugged his data into the quantum theory, he found that electrons do not travel in nice, neat orbits (Bohr model) • Instead, found that electrons travel in 3D regions around the nucleus

QUANTUM MECHANICS • Schrodinger • Schrodinger’s equation is used to find the greatest probable location/area of the Hydrogen atom electron (in the ground state)

QUANTUM MECHANICS • Quantum Theory • Ground State-normal location of an electron • Excited State-one ring up from the normal location • When excited electron falls back to the ground state, a photon is given off • Energy of the photn is equal to the difference in energy between the excited state and ground state • Hydrogen gives off specific colors because its electrons move from ring 2 to ring 1; Neon gives off a different color because its electrons move from ring 3 to ring 2

LIGHT AND ELECTRONS REVIEW • Light was first thought to be wavelike • Equation for the speed of light is c =  • Photoelectric effect challenges this because only certain frequencies of light could knock off electrons • Max Planck’s experiment proved that light could be a particle • Einstein’s dual wave particle theory says that light is ACTUALLY small particles (photons) that move in wave like patterns • Equation for energy of a photon is E = h • Bohr found that electrons orbit the nucleus in specific orbitals/energy levels

LIGHT AND ELECTRONS REVIEW • Electrons as Waves: • 1924 – Louis de Broglie asked “Could electrons have a dual wave particle nature like light?” • Similarities between waves and electrons • Waves have specific frequencies and electrons have specific orbits/energy levels • Waves and electrons can both be bent (diffraction) • Waves and electrons can both overlap and interfere with each other (interference) • DeBroglie’s equation combines Einstein and Planck’s equations • mv/ = h • (Anything with mass and velocity has a wavelength, so electrons have wavelengths)

QUANTUM NUMBERS and ATOMIC ORBITALS • REVIEW • Energy levelsSpecific energies electrons can have • Quantum of energyamount of energy required to move an electron from one energy level to another energy level • The amount of energy an electron gains or loses in an atom is not always the same • Energy levels in an atom are not equally spaced • Higher energy levels are closer together • Modern description of the electrons in atoms, quantum mechanical model, comes from the mathematical solutions to the Schrodinger equation • The quantum mechanical model determines the allowed energies an electron can have and how likely it is to find the electron in various locations

QUANTUM NUMBERS and ATOMIC ORBITALS • QUANTUM NUMBERS • Quantum numbers are used to describe the location and behavior of an electron (zip code for electrons) • First quantum number = Principal = n • Second quantum number = Angular Momentum • Third quantum number = Magnetic Quantum • Fourth quantum number = Spin Quantum

QUANTUM NUMBERS and ATOMIC ORBITALS • Principal (first) quantum number = n • Main quantum number • Describes the size of the electron cloud (the smaller the number, the smaller the cloud) • ALSO, shows the distance from the nucleus, the smaller the number, the closer the cloud is to the nucleus • Called energy levels or shells • Positive integers (1,2,3,4,…) • Symbol is n • Each energy level has a maximum number of electrons it can hold n# Electrons 1 2 2 8 3 18 4 32 • Example: Energy level 1 2 electrons close to the nucleus small electron cloud

QUANTUM NUMBERS and ATOMIC ORBITALS • Second Quantum Number: • Each energy level has sublevels • The number of sublevels is equal to n • Example: Energy level 1 has 1 sublevel • Sublevels are called: s,p,d,f

QUANTUM NUMBERS and ATOMIC ORBITALS • Third Quantum Number • Divides sublevels into orbitals • Also tells the shape the electron is moving in • The number of orbitals for each level is: • S has 1 • P has 3 • D has 5 • F has 7 • The number of orbitals for an energy level is equal to n2 • Example: 2nd Energy level • n2 = 4 • 1s, 3p • Each orbital can only hold a maximum of 2 electrons • Shapes of orbitals: • S is spherical • P is peanut shaped • D is daisy shaped • F is unknown

QUANTUM NUMBERS and ATOMIC ORBITALS • Fourth Quantum Number: • Describes the electron spin • Both electrons in an orbital are negative, so they repel each other and spin in opposite directions • Use arrows to represent electrons

QUANTUM NUMBERS and ATOMIC ORBITALS • Pauli Exclusion Principle: • No two electrons in an atom can have the same set of 4 quantum numbers because electrons repel each other • Example: 2 electrons may both be: in the first energy level (same first number) sitting in an s sublevel (same second # moving in a sphere shape (same third #) BUT one electron spins clockwise and one spins counter clockwise ( which means they have different fourth #s)

ELECTRON CONFIGURATIONS • Example 1: Map out the quantum numbers for all the electrons in Hydrogen • Find the # of electrons in hydrogen (atomic number will give you this number) • What order do you fill in s, p, d, f in the rings?

ELECTRON CONFIGURATIONS • Diagonal RulePattern that shows the order the electrons fill in the orbitals: Some People Do Forget 1s 2s 2p 3s 3p 3d 4s 4p 4d 4f 5s 5p 5d 5f 6s 6p 6d 7s 7p Notice that the electrons do not fill in all of the level 3 first (3s, 3p, 3d) and then move to level 4 Instead, electrons fill in the orbitals in the order that is easiest to them (easier for an electron to fill in a 4s before it fills in a 3d) Aufbau Principle: Electrons have to fill in the lowest (easiest) energy level or orbital first

Hund’s Rule: Every orbital must get one electron first, before you double up. “Cookie Rule” Example 2: Helium ELECTRON CONFIGURATIONS

Example 3: Lithium Example 4: Fluroine ELECTRON CONFIGURATIONS

Orbital Notation drawing out configurations with arrows Electron Configuration Notation/Superscript Notation: writing configurations with superscripts to represent electrons ELECTRON CONFIGURATIONS

ELECTRON CONFIGURATIONS • Do Orbital Notation and Electron Configuration for the following: • Zn • I • Cl • Mg • As

NOBLE GAS CONFIGURATIONS • Noble Gas Configurations: • Write out the superscript notations for: Neon: Sulfur: Sulfur has the same configuration as Neon plus a 3s23p4 So, you could use the noble gas as a shortcut and write Sulfur’s configuration as [Ne] 3s23p4 OR [Ne] • Noble gas configuration: write the noble gas (group 18) that comes directly before the element in question and then add the rest of the configuration • Practice: • Write the noble gas superscript notation for the following elements. C Np W Sn

DOT DIAGRAMS • Lewis Dot Diagrams: • Way to show the number and position of electrons on the outermost energy level • Since the energy levels all overlap and cover one another, only the outermost energy level is able to bond with other elements • The electrons involved in bonding are called the valence electrons (to get these electrons look at the column number)

DOT DIAGRAMS • Lewis Dot Diagrams: • Chemical symbol + Number of valence electrons • The rules for orbitals still apply, so no side can have more than two dots, and each “p” orbital side gets one dot, before you double up X p1 s orbital s p orbitals p2 p3

DOT DIAGRAMS • Noble gases have a full valence • There are no empty spaces so the element does not need any more electrons • Stable octet – 8 electrons in the valence so the element does not want to bond (this means it is stable) • Only the noble gases have a stable octet

DOT DIAGRAMS • Practice: Write the noble gas superscript notation and then draw the dot diagram for the following: • V • Br • Al • K

LIGHT and ELECTRONS