Unit 10 Electrons in Atoms

Light as Energy Wave Theory of Light: James Clerk Maxwell (mid-1800s) Electromagnetic radiation – a form of energy that exhibits wavelike behavior as it travels through space Visible light – a form of electromagnetic radiation that is perceivable to human beings and is seen in the colors of the rainbow (ROY G. BIV)

Wave Diagram

Wave Vocab: Crest – the top of a wave Trough – the bottom of a wave Wavelength – the distance from crest to crest or trough to trough in a wave Frequency – the number of wavelengths that pass a given point in a set amount of time Origin – the center/start or midpoint of a wave Amplitude – the distance from the origin to the crest or the trough of a wave Speed of light – c – the rate at which all forms of electromagnetic radiation travel through a vacuum = 3.00 x10 8 m/s

Wave Theory of Light Light as a Wave A. Wavelength (  “lambda”) Distance between corresponding parts on a wave (m) B. Frequency (  “nu”) # of peaks that pass a given point each second (cycles/sec  Hz)

Wave Theory of Light B. Frequency (  “nu”) # of peaks that pass a given point each second (cycles/sec  Hz)

Comparing Waves High frequency = Short Wavelength Low frequency = Long Wavelength

Wave Theory of Light C. Wave Velocity (speed) = c Distance that a peak travels in a unit of time (generally one second) Units: m/s Speed of light = c = 3.00 x 10 8 m/s D. Amplitude The height of a wave (measured in meters)

Wave Equation: Wavelength = speed of light / frequency Wavelength and frequency are inversely proportional

c = Example #1: The wavelength of the radiation which produces the yellow color of sodium vapor light is 5.89 x m. What is the frequency of this radiation?

c = Example #2: What is the wavelength of a microwave having a frequency of 3.44 x 10 9 Hz?

Visible Light Consists of all the colors of the rainbow ROY G. BIV Can be separated by a prism

Electromagnetic Spectrum All forms of light both visible and non-visible Ranges from radio waves to gamma rays

Particle Theory of Light First studied by Max Planck around 1900 Came about to try and explain the parts of light behavior that didn’t fit Maxwell’s wave theory

Evidence for Particle Theory Photoelectric Effect: - Light is emitted in small discrete amounts called “photons” Photon: a single “particle” of light

Quantum Concept When an electron absorbs a photon of energy, the electron jumps from the ground state to its excited state Ground state – lowest energy level an electron occupies Excited state – temporary state when an e- is at a higher energy level

Quantum Concept quantum: the minimum amount of energy that can be gained or lost by an atom Photon: a single “particle” of light Planck’s Constant = h = x J. s

Comparing Energy and Frequency Energy and frequency are directly proportional They both increase or decrease together SO…. Energy and wavelength are inversely proportional

E = h  Example #1: What is the energy of a photon from the violet portion of the rainbow if it has a frequency of 7.23 x s -1 ?

E = h  Example #2: What is the energy of a photon from the green portion of the rainbow if it has a wavelength of 4.90 x m?

Dual Theory of Light Albert Einstein – 1905 Light is both a wave and a particle!!!! Demonstrations: Light bends

Atomic Emission Spectra The set of frequencies of light that are emitted when atoms of a gas are energized with electricity Each atom has its own distinct AES Intensive Property Hydrogen Atom Spectrum

Quantum Theory and the Atom Neils Bohr extended his “Planetary Model” of the atom to address light Studied the Hydrogen atom and how different light was emitted when the atoms electrons moved between energy levels

Hydrogen’s Line Spectrum Bohr stated that electrons exist in set pathways called energy levels Energy levels are represented by the letter n So, the first energy level is n = 1,etc. The first energy level is closest to the nucleus, etc.

Hydrogen’s Line Spectrum

What determines the color of light? Bohr concluded that the color of light depends on the fall of the electron from the excited state to the ground state Balmer, Lyman and Paschen each studied specific energy levels and the light emitted when those energy levels are involved

Heisenberg’s Uncertainty Principle Both the velocity and position of an electron cannot be determined simultaneously You can determine either the speed OR the location at a given time but not both Since an electron is so small and the methods for determining position and velocity can’t be done simultaneously, doing one creates error in the other

Quantum Mechanical Model of the Atom Erwin Schrodinger – 1926 Explained electron movement as wavelike rather than set as in the Bohr model (remember from Chapter 4) Gave electrons regional locations called orbitals and based them on probability calculations

Schrodinger’s Atom

Electron Arrangement in the Electron Cloud Electron locations are designated by a system of letters and numbers called “quantum numbers” These numbers give you the location of an electron ranging from the most general location to the most specific location

1. Principle Energy Level (n) The most general of the information Whole numbers from 1,2,3,4, etc. Remember 1 is closest to the nucleus, 2 next, etc. The principal energy level is the same as the row number of the periodic table Example: Valence electrons in calcium reside in the n = 4 energy level

2. Sublevel One step more specific than the energy level Letters s, p, d, and f Energy level 1 has only sublevel s Energy level 2 has s and p Energy level 3 has s, p, and d Energy level 4-7 have s, p, d, and f

3. Orbital The most specific piece of information is about the number and location of the electrons within the sublevel Can be represented by an exponent or arrows depending on the method of notation The s sublevel has 1 orbital The p sublevel has 3 orbitals The d sublevel has 5 orbitals The f sublevel has 7 orbitals Orbital - region within a sublevel where an e - can be found (homes for e - )

Rules for how the electrons fill into the electron cloud: Aufbau Principle: electrons fill from the lowest energy level to the highest (they don’t skip around) Pauli Exclusion Principle: each orbital can hold a maximum of 2 electrons at a time (they must have opposite spins) Hund’s Rule: orbitals of equal energy in a sublevel must all have 1 electron before the electrons start pairing up

Electron Configurations The system of numbers and letters that designates the location of the electrons 3 major methods: Full electron configurations Abbreviated/Noble Gas configurations Orbital diagram configurations

Full Electron Configuration Example Notation: 1s 2 2s 1 (Pronounced “one-s-two, two-s-one”) A. What does the coefficient mean? Principle energy level B. What does the letter mean? Type of orbital (sublevel) C. What does the exponent mean? # of electrons in that orbital

Full Electron Configuration 1.Determine the total number of electrons the atom has (for neutral atoms it is equal to the atomic number for the element). Example: F atomic # = # of p + = # of e - = 2.Fill orbitals in order of increasing energy. 3.Make sure the total number of electrons in the electron configuration equals the atomic number.

Aufbau Chart (Order of Energy Levels) When writing electron configurations: d sublevels are n – 1 from the row they appear in f sublevels are n – 2 from the row they appear in

Full Electron Configuration Practice Write electron configurations for the following elements: He: P: Rh: Br: Ca: Ce:

Abbreviated/Noble Gas Configuration i.Where are the noble gases on the periodic table? ii.Why are the noble gases special? iii. How can we use noble gases to shorten regular electron configurations?

Abbreviated/Noble Gas Configuration Example: Barium 1.Look at the periodic table and find the noble gas in the row above where the element is. 2.Start the configuration with the symbol for that noble gas in brackets, followed by the rest of the electron configuration.

Abbreviated/Noble Gas Configuration Practice! Write Noble Gas Configurations for the following elements: Rubidium: Bismuth: Arsenic: Zirconium:

Quantum Numbers: Hotel Analogy Recall there were three values we talked about which indicated the location of an electron: Principal energy level: n (row of the periodic table) Sublevel: s, p, d, f Orbital: each orbital can hold up to 2 e - These two electrons form an “electron pair” which must have opposite spin (point in opposite directions)

Quantum Numbers: Hotel Analogy Principal energy level: n (= row of periodic table) Sublevel: s, p, d, f Orbital: each orbital can hold 2 e - Think of these in terms of a hotel: Floor – most general; hotels have many floors (like the energy level in an atom) Wing – each floor has a few wings or corridors (like the sublevel in an atom) Room – most specific location (like the orbital in an atom)

Orbital Diagrams Orbital diagrams use boxes (sometimes circles) to represent energy levels and orbitals. Arrows are used to represent the electrons. = orbital sublevels

Orbital Diagrams Don’t forget - orbitals have a capacity of two electrons!! Two electrons in the same orbital must have opposite spin so draw the arrows pointing in opposite directions. Example: oxygen1s 2 2s 2 2p 4 1s 2s 2p Increasing Energy 

Drawing Orbital Diagrams 1. First, determine the electron configuration for the element. 2. Next draw boxes for each of the orbitals present in the electron configuration. Boxes should be drawn in order of increasing energy (see the Aufbau chart). 3. Arrows are drawn in the boxes starting from the lowest energy sublevel and working up. This is known as the Aufbau principle. Add electrons one at a time to each orbital in a sublevel before pairing them up (Hund’s rule) The first arrow in an orbital should point up; the second arrow should point down (Pauli exclusion principle) 4. Double check your work to make sure the number of arrows in your diagram is equal to the total number of electrons in the atom. # of electrons = atomic number for an atom

Drawing Orbital Diagrams Practice! Draw orbital diagrams for the following elements: Sulfur: Nickel:

Drawing Orbital Diagrams (Answers) Sulfur: Nickel:

Orbital Diagrams: Valence e - & Ionic Charge If the orbital diagram for sulfur looks like this: How many unpaired electrons are in the valence shell of a sulfur atom? Does this make sense regarding the charge on a sulfur ion?