Electronic Structure and the Periodic Table Unit 6 Honors Chemistry

Wave Theory of Light James Clerk Maxwell Electromagnetic waves – 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 (  “lambda”) the distance from crest to crest or trough to trough in a wave Units: m, nm (1 m = 10 9 nm) Frequency (  “nu”) the number of wavelengths that pass a given point in a set amount of time (generally in 1 second) Units: Hertz (Hz), 1/s, or s -1

Wave Vocab: Amplitude – the distance from the origin to the crest or the trough of a wave Height (or intensity/brightness) of 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

Comparing Waves As Wavelength increases, frequency _______________. As Wavelength decreases, frequency _______________.

Wave Equation One equation relates speed, frequency and wavelength: c =

Example The wavelength of the radiation which produces the yellow color of sodium vapor light is nm. What is the frequency of this radiation? c =

The Electromagnetic Spectrum Complete range of wavelengths and frequencies Mostly invisible

What is Color? TED Ed Video: What is color?

The Visible Spectrum Continuous spectrum: components of white light split into its colors, ROY G. BIV From 390 nm (violet) to 760 nm (red) Can be split by a prism

How do we see color? TED Ed Video: How we see color

Max Planck – Particle Theory of Light Light is generated as a stream of light particles called PHOTONS Equation: E = h h =Plank’s constant= x J · s)

Example #1 (a) If the frequency of a ray of light is 5.09 x Hz, calculate the energy, in joules, of a photon emitted by an excited sodium atom. (b) Calculate the energy, in kilojoules, of a mole of excited sodium atoms.

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?

Bohr Model of the Atom When an electron absorbs a photon of energy, the electron jumps from the ground state to an excited state Ground state – lowest energy level an electron occupies Excited state – temporary state when an electron is at a higher energy level

Line Spectra Pattern of lines produced by light emitted by excited atoms of an element Unique for every element Used to identify unknown elements

Explanation of Line Spectra Niels Bohr Energy of an electron is quantized: can only have specific values. Energy is proportional to energy level.

Explanation of Line Spectra Electron will drop from excited state to ground state and will emit energy as a photon during the fall. Video: Atomic Emission Animation

Photoelectric Effect – Nobel Prize in Physics 1921 to Einstein Occurs when light strikes the surface of a metal and electrons are ejected. Practical uses: Automatic door openers Ted Ed Video: Is Light Actually a Wave or Particle?

Conclusion… Light not only has wave properties but also has particle properties. These massless particles, called photons, are packets of energy. Light has a dual nature!

Quantum Mechanics Quantum mechanics: atomic structure based on wave- like properties of the electron Schrödinger: wave equation that describes hydrogen atom

Heisenberg Uncertainty Principle The exact location and speed of an electron cannot be determined simultaneously (if we try to observe it, we interfere with the particle) You can know either the location or the velocity but not both Electrons exist in electron clouds and not on specific rings or orbits like in the Bohr model of the atom

Quantum Numbers Quantum numbers – a system of four numbers used to represent the most probable location of an electron in an atom They range from the most general locator to the most specific Analogy... state = energy level, n city = sublevel, l address = orbital, m l house number = spin, m s

1. Energy Level Principal Quantum Number: n Always a positive integer (1, 2, 3,…7) Indicates size of orbital, or how far electron is from nucleus Larger n value = larger orbital or farther distance from nucleus Similar to Bohr’s energy levels or shells

n = 1 n = 2 n = 3 n = 4 n = 5 n = 6 n = 7 n = row number on periodic table for a given element n in relation to the Periodic Table

Indicates shape of orbital 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 2. Sublevel Angular Momentum Quantum Number: l

In order of increasing energy the sublevels generally go: s < p < d < f HOWEVER, there is some overlapping of sublevels at higher energy levels Ex.) 4s vs. 3d

3. Orbital The most specific piece of information is about the number and location of the electrons within the sublevel 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 - ) Every orbital can hold 2 electrons!

Orbitals Orbital = electron containing area (houses for electrons) No more than 2 e- assigned to an orbital Orbitals grouped in s, p, d (and f) subshells

Shapes of Atomic Orbitals s = spherical p = peanut d = dumbbell (clover) f = flower

Capacities of levels, sublevels, and orbitals Principal Energy level (n) Sublevels Present (s, p, d, or f) Number of Orbitals Present s p d f Total Number of Orbitals Maximum Number of Electrons in Energy Level

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 (and 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

Why are these incorrect?

Electron Configuration Definition: describes the distribution of electrons among the various orbitals in the atom Represents the most probable location of the electron! EOS

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 sublevel – s, p, d, or f C. What does the exponent mean? # of electrons in that sublevel

Steps to Writing Full Electron Configurations 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 (see Aufbau Chart). 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

In order of increasing energy the sublevels generally go: s < p < d < f HOWEVER, there is some overlapping of sublevels at higher energy levels Ex.) 4s vs. 3d

Writing Electron Configurations Nitrogen: Helium: Phosphorous: Rhodium: Bromine: Cerium:

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: Tin 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: Sufur: Rubidium: Bismuth: Zirconium:

Orbital Diagrams Another way of writing configurations is called an orbital diagram. (also called orbital notation) Another way of writing configurations is called an orbital diagram. (also called orbital notation) One electron has n = 1, l = 0, m l = 0, m s = + ½ Other electron has n = 1, l = 0, m l = 0, m s = - ½

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 spins 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

Orbital Configurations for Nitrogen Full Electron Configuration: Orbital Diagram:

Orbital Configurations for Nickel Full Electron Configuration: Orbital Diagram:

Exceptions to the Filling Order Rule (Cr, Cu)—these will not be on test!

Valence Electrons Definition: Electrons in the outermost energy levels They determine the chemical properties of an element! ***Write the noble gas configuration...the valence electrons are the ones beyond the noble gas core in the highest energy level

Valence Electrons and Core Configuration (Shorthand) What is the shorthand notation for S? EOS Sulfur has six valence electrons

Configurations of Ions Cations: Formed when metals lose e – in highest principal energy level. Example: (Z = 11) Na EOS (Z = 11) Na +

Configurations of Ions Anions: Formed when non-metals gain e – to complete the p sublevel EOS -

Transition Metals Transition metals (and p block metals) lose e – from the highest principal energy level (n) FIRST, then lose their d electrons! EOS Zr = [Kr] 5s 2 4d 2 Zr +2 = [Kr] 4d 2

Periodic Trends!

Periodic Properties & Trends Electronegativity –Ability of an atom to pull e - towards itself –Linus Pauling: developed scale to demonstrate different electronegativity strengths –Increases going up and to the right Across a period  more protons in nucleus = more positive charge to pull electrons closer Down a group  more electrons to hold onto = element can’t pull e - as closely

Electronegativity –Ability of an atom to pull e - towards itself –Across a period  more protons in nucleus = more positive charge to pull electrons closer –Down a group  more electrons to hold onto = protons in nucleus can’t pull e - as closely Periodic Properties & Trends

Atomic Radius –Distance between the nucleus and the furthest electron in the valence shell –Increases going down and to the left Down a group  more energy shells = larger radius Across a period  elements on the right can pull e - closer to the nucleus (more electronegative) = smaller radius *Remember* –LLLL  Lower, Left, Large, Loose

Periodic Properties & Trends Atomic Radius –Increases going down and to the left *Remember* LLLL  Lower, Left, Large, Loose

Memory Device LLLL: Lower Left, Larger Atoms

Periodic Properties & Trends Ionic Radius –Radius of an atom when e - are lost or gained  different from atomic radius –Ionic Radius of Cations Decreases when e- are removed –Ionic Radius of Anions Increases when e- are added

Sizes of Ions CATIONS are SMALLER than the atoms from which they are formed. Size decreases due to increasing he electron/proton attraction. Li,152 pm 3e and 3p Li +, 78 pm 2e and 3 p +

Sizes of Ions ANIONS are LARGER than the atoms from which they are formed. Size increases due to more electrons in shell. F, 71 pm 9e and 9p F -, 133 pm 10 e and 9 p -

Trends in Ion Sizes Active Figure 8.15 Trends in ion sizes are the same as atom sizes.

Periodic Properties & Trends Ionization Energy –Energy required to remove an e- from the ground state –1 st I.E. = removing 1 e -, easiest –2 nd I.E. = removing 2 e -, more difficult –3 rd I.E. = removing 3 e -, even more difficult Ex.) B --> B + + e- I.E. = 801 kJ/mol Ex.) B + --> B +2 + e- I.E.2 = 2427 kJ/mol Ex.) B +2 --> B +3 + e- I.E.3 = 3660 kJ/mol

Periodic Properties & Trends Ionization Energy Increases going up and to the right –Down a group  more e - for the nucleus to keep track of = easier to rip an e - off –Across a period  elements on the right can hold electrons closer (more electronegative) = harder to rip an e - off

Memory Device LLLL: Lower Left, Larger Atoms; Looser electrons LLLL: Lower Left, Larger Atoms; Looser electrons

Periodic Properties & Trends Metallic Character –How “metal-like” an element is Metals lose e - –Most Metallic: Cs, Fr –Least: F, O –Increases going down and to the left Think about where the metals & nonmetals are located on the periodic table to help you remember!

Electron Affinity Electron affinity is the energy involved when an atom gains an electron to form an anion. Some elements GAIN electrons to form anions. A(g) + e- ---> A - (g) E.A. = ∆E

Trends in Electron Affinity Trend in a group: Affinity for e - increases going up a group Trend in a series or period: Affinity for e - increases going across a period to the right

Electron Affinity Note that the trend for E.A. is the SAME as for I.E. !

A Summary of Periodic Trends Remember LLLL!!