6.1-6.3

6.4-6.7

Name: ____________________________ Hr: ____

6.8-6.9

6  Electronic Structure of Atoms 6.1 The Wave Nature of Light I can…

Objectives 6.5 Quantum Mechanics and Orbitals I can…

 Describe the wave properties and characteristics

 Explain the concepts of orbital, electron density,

of electromagnetic radiation.

 Calculate , , or E of any wave given one of the other quantities.

 Qualitatively relate the size of wavelength to size of frequency and to size of energy.

 State the seven types of electromagnetic radiation (EMR) in order of energy, frequency, and wavelength.

and probability.

 Explain what probability waves are and explain 2 the significance of ψ .

 Describe electrons by level, sublevel (orbital shape), orientation (which orbital within the sublevel) and spin. 6.6 Representations of Orbitals I can…

6.2 Quantized Energy and Photons I can…

 Draw the general shapes of orbitals (the

 Use Planck’s equation to determine the energy of

 State the relationship between n and the average

photons of radiation at different frequencies.

 Describe the photoelectric effect.  Explain how the photoelectric effect provides evidence that light (waves) must be particles (photons).

standing waves of an electron). distance from the nucleus. 6.7 Many-Electron Atoms I can…

 Explain why electrons with the same level but different sublevel possess different energies.

6.3 Line Spectra and the Bohr Model I can…

 Explain the concepts of electron spin as it

 Describe the differences among a continuous

 Explain the Pauli exclusion principal and its

emission spectrum, a bright line spectrum, and an absorption spectrum.

 List the assumptions made by Bohr in his model of the hydrogen atom.

 Explain the concept of an allowed energy state and how this concept is applied to quantum theory.

 Define the terms ground and excited state. 6.4 The Wave Behavior of Matter I can…

 Describe the uncertainty principle and explain the limitation it places on our ability to define the location and momentum of an electron.

relates to quantum theory. implications for electrons in an orbital (and electron configurations). 6.8 Electron Configurations I can…

 State the three rules and give examples of how they govern electron assignment: o the Aufbau Principle o the Pauli Exclusion Principle o Hund’s Rule

 Draw the energy diagram for multi-electron atoms (draw the orbital diagram).

 Write the electron configuration, condensed electron configuration (noble gas), and electron dot notation for an element.

6.9 Electron Configuration and Periodic Table I can…

 Show how the regions of the periodic table correspond to the orbitals; identify the s-block, pblock, d-block, and f-block.

 Write the valence electron configuration for any element based upon its place in the periodic table.

Chapter 6 Notes – Electronic Structure of Atoms 6.1 The Wave Nature of Light  The electronic structure of an atom refers to the arrangement of _____________. o To understand the electronic structure of atoms, one must understand the nature of electromagnetic radiation.  Electromagnetic radiation is characterized by its wave nature o All waves have a characteristic wavelength, lambda), and amplitude, A.  The units of  are in meters (m), nanometers (10-9), and angstroms (10-10) o The frequency, nu), of a wave is the number of cycles which pass a point in one second.  The units of  are Hertz (1 Hz = 1 s–1) or MHz (106 Hz). o For waves traveling at the same velocity, the longer the wavelength, the ____________ the frequency.  All electromagnetic radiation travels at the same velocity: the speed of light (c), 3.0  108 m/s. o For light, speed, c = ,  The electromagnetic spectrum is a display of the various types of electromagnetic radiation arranged in order of decreasing wavelength (_______________ frequency and energy). o Radio, Microwaves, Infrared, Visible Light (ROYGBIV), Ultraviolet, X rays, and Gamma rays

Speed of Light Problem What is the wavelength of a microwave having a frequency of 3.44 x 109 Hz?

c   

Speed of Light Problem 2 A helium-neon laser emits light with a wavelength of 633 nm. What is the frequency of this light?  1 10 -9 m   633 nm    1 nm 

c  



6.2 Quantized Eenergy and Photons  Some phenomena can't be explained using a wave model of light: o Blackbody radiation: emission of light from hot objects. o The photoelectric effect: emission of electrons from metal surfaces on which light shines. o Emission spectra: emission of light from electronically excited gas atoms.

Blackbody Radiation  Planck investigated black body radiation. o He proposed that energy can only be absorbed or released from atoms in certain amounts. o These amounts are called quanta.  A quantum is the smallest amount of energy that can be emitted or absorbed as electromagnetic radiation. o The relationship between energy and frequency is:  E = _______ o where h is Planck’s constant (6.63 x 10–34 J·s).  To understand quantization consider the notes produced by a violin (continuous) and a piano (quantized): o A violin can produce any note when the fingers are placed at an appropriate spot on the finger board. o A piano can only produce notes corresponding to the keys on the keyboard.

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Chapter 6

Photon Energy Problem What is the energy of a photon from the violet portion of the rainbow if it has a frequency of 7.23 x 1014 Hz?

E  h E = (6.63 × 10-34 J·s)(7.23 × 1014 1/s) E=

Photoelectric Effect  Einstein assumed that light traveled in energy packets called photons. o The energy of one photon, E = h. o Light shining on the surface of a metal can cause electrons to be ejected from the metal.  The electrons will only be ejected if the photons have sufficient energy: o Below the threshold frequency ____ electrons are ejected. o Above the threshold frequency, the excess energy appears as the ________ energy of the ejected electrons.

Energy and Wavelength Combined  Combing the speed of light equation and Plank’s equation yields: E

6.3 Line Spectra and the Bohr Model Line Spectra  Radiation composed of only one wavelength is called _________________.  Radiation that spans a whole array of different wavelengths is called ________________.  When radiation from a light source such as a light bulb is separated into its different wavelength components, a spectrum is produced.  White light can be separated into a continuous spectrum of colors. o A rainbow is a continuous spectrum of light produced by dispersal of sunlight by raindrops or mist. o Note that there are no dark spots on the continuous spectrum which would correspond to different lines.  Not all radiation is continuous.  A gas placed in a partially evacuated tube and subjected to a high voltage produces single colors of light.  The spectrum that we see contains radiation of only specific wavelengths; this is called a line spectrum.

Bohr’s Model    

Rutherford assumed the electrons orbited the nucleus analogous to planets around the sun. However, a charged particle moving in a circular path should lose energy. This means that the atom should be unstable according to Rutherford’s theory. Bohr noted the line spectra of certain elements and assumed the electrons were confined to specific energy states. These were called orbits.  Bohr model is based on three postulates: 1. Only orbits of specific radii, corresponding to certain definite energies, are permitted for electrons in an atom. 2. An electron in a permitted orbit has a specific energy and is an "allowed" energy state. 3. Energy is only emitted or absorbed by an electron as it moves from one allowed energy state to another. 4. The energy is gained or lost as a photon.

The Energy States of the Hydrogen Atom

 The first orbit in the Bohr model has n = 1 and is closest to the nucleus.  Electrons in the Bohr model can only move between orbits by absorbing and emitting energy in quanta (E = h.

 The ground state = the lowest energy state.  An electron in a higher energy state is said to be in an excited state.

Limitations of the Bohr Model  The Bohr Model has several limitations: o It cannot explain the spectra of atoms other than hydrogen. o Electrons do not move about the nucleus in ______________ orbits.

Electronic Structure of Atoms

 However the model introduces two important ideas: o The energy of an electron is ________________: electrons exist only in certain energy levels described by quantum numbers. o Energy gain or loss is involved in _____________ an electron from one energy level to another

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6.4 The Wave Behavior of Matter  Knowing that light has a particle nature, it seems reasonable to ask whether matter has a wave nature.  This question was answered by Louis deBroglie.  Using Einstein’s and Planck’s equations, deBroglie derived:

  The momentum, mv, is a particle property, whereas  is a wave property.  Matter waves is the term used to describe wave characteristics of material particles. o Therefore, in one equation de Broglie summarized the concepts of waves and particles as they apply to low-mass, high-speed objects. o As a consequence of deBroglie’s discovery, we now have techniques such as X-ray diffraction and electron microscopy to study small objects.

The Uncertainty Principle  Heisenberg’s uncertainty principle: You can’t precisely know both the position and momentum of a particle at the same time.

x   In many cases, our uncertainty of the whereabouts of an electron is greater than the size of the atom itself!

6.5 Quantum Mechanics and Atomic Orbitals  Schrödinger proposed an equation containing both wave and particle terms. o Solving the equation leads to wave functions, .  The square of the wave function, , gives the __________________ of finding an electron in a cetain location. o  is called the probability density or electron density. o The “electron ___________” is another way of expressing the probability density.  A region of high electron density is one where there is a high probability of finding an electron.  The wave function gives the ___________ of the electron’s orbital.

Orbitals and Quantum Numbers    

Solving the wave equation gives a set of wave functions, or orbitals, and their corresponding energies. Each orbital describes a spatial distribution of electron density. An ____________ is described by a set of three quantum numbers. While we no longer need to know the symbols and numbers associated with these quantum numbers for the AP Test, we still need to understand the concepts related to them to understand and explain the electronic structure of atoms.  The three “quantum numbers” used to describe orbitals are: 1. Principle quantum number (n) – energy level 2. Sublevel – orbital shape (s, p, d, f) 3. Orientation – which orbital on the sublevel

Principal quantum number (Energy Level)

 The principal quantum number, n, describes the energy level on which the orbital resides. o Describes distance from the nucleus and general ___________. o n = 1, 2, 3, 4, …. o The higher the energy level the larger the atom and the _____________ the electron is from the nucleus.

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Chapter 6

Sublevel (Shape)  Each energy level contains _____________ o The number of sublevels on a level is equal to the energy level (n).  1st energy level has ___ sublevel  2nd energy level has ___ sublevels  Each sublevel contains orbitals. o Orbital: a three-dimensional region around the nucleus in which an electron moves and is found ______ of the time.  Each orbital can hold up to two electrons. o The total number of orbitals on a level = _____. o Each sublevel has a different shape of orbital on the level.  These shapes are represented by the symbols s, p, d, or f.

Orientation

 Each shape of orbital has a different number of orientations in which they can exist on a sublevel. o s → ___ o p → ___ o d → ___ o f → ___

Other Related Terms for the AP Test

 Orbitals with the same value of n form an electron __________.  Different orbital types within a shell are _____________.

6.6 Representation of Orbitals s Orbitals  Each level has _____ s shaped (spherical) sublevel o n = 1, 2, 3 …  Only ___ orientation on sublevel o An s sublevel can hold ___ electrons  As n increases, the s orbitals get __________. o This true of all orbitals!

p Orbitals  ___ energy level and above have a p sublevel o n = 2,3,4,…  ___ orientations on each sublevel o They have two lobes, but are still only 1orbital!  p sublevels can hold up to ___ electrons

d Orbitals  ___ energy level and above have a d sublevel o n = 3,4,5,…  ___ orientations on each sublevel  d sublevels can hold up to ____ electrons

f Orbitals  ___ energy level and above have a f sublevel o n = 4,5,6,…  ___ orientations on each sublevel  f sublevel can hold up to ____ electrons

Electronic Structure of Atoms

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6.7 Many-Electron Atoms Orbitals and Their Energies  For a one-electron hydrogen atom, orbitals on the same energy level have the same energy. o That is, they are _______________.  As the number of electrons increases, though, so does the repulsion between them. o Therefore, in many-electron atoms, orbitals on the same energy level are no longer degenerate. o Here, the sublevels are degenerate.

One electron atom

Many-electron atom

Electron Spin and the Pauli Exclusion Principle  In the 1920s, it was discovered that two electrons in the same orbital do not have exactly the same energy.  The “_____” of an electron describes its magnetic field, which affects its energy  This led to a fourth quantum number, the spin quantum number, ms. (_________) o The spin quantum number has only 2 allowed values: ____ and ____.  Pauli’s exclusion principle states that no two electrons can have the _______ set of 4 quantum numbers. o Therefore, two electrons in the same orbital must have ____________ spins.

6.8 Electron Configurations

 Electron configurations tell us how the electrons are distributed among the various orbitals of an atom.

 The most stable configuration or ground state is that in which the electrons are in the ___________ possible energy state.  When writing ground-state electronic configurations: 1. Electrons fill orbitals in order of increasing energy with no more than two electrons per orbital. (__________) 2. No two electrons can fill one orbital with the same spin (__________). 3. For degenerate orbitals, electrons fill each orbital singly before any orbital gets a second electron. (________)  How do we show spin? o An arrow pointing upwards has ms = + ½ (spin up). o An arrow pointing downwards has ms = – ½ (spin down).

Hund's Rule  Hund’s rule: For degenerate orbitals, the lowest energy is attained when the number of electrons with the same spin is maximized. o Thus electrons fill each orbital singly with their spins parallel before any orbital gets a ____________ electron. o By placing electrons in different orbitals, electron-electron _______________ are minimized.

Condensed Electron Configurations  Condensed, Shorthand Notation, or Noble Gas o Like electron configuration except the inner level electrons are described by writing the last noble gas in brackets and then describing the other electrons.  lead 1s22s22p63s23p64s23d104p65s24d105p66s24f145d106p2  [Xe]_____________

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Chapter 6

6.9 Electron Configurations and the Periodic Table  We fill orbitals in increasing order of energy.  Different blocks on the periodic table correspond to different types of orbitals.  The period number is the value of ___. o Groups 1A and 2A have their s orbitals being filled. o Groups 3A – 8A have their p orbitals being filled.  The s-block and p-block of the periodic table contain the ______________ elements or main-group elements.  d-block o Groups 3-12 have their d orbitals being filled.  f-block o The lanthanides and actinides have their f orbitals being filled.  Note that the 3d orbitals fill after the 4s orbital. Similarly, the 4f orbitals fill after the 5d orbitals.

Electron Configurations of Ions  When an ion is formed, electrons are removed from the highest energy levels first (farthest from the nucleus). o Na d 1s22s22p63s1 o Na+ d ____________ o Sn d [Kr]5s24d105p2 o Sn4+ d ____________  For transition metals, first remove the highest level (n)p electrons,then (n)s electrons, then the (n-1)d electrons until the cation charge is reached. o Fe d [Ar]4s23d6 o Fe2+ d ________ o Fe3+ d ________

Are atoms weakly attracted to a magnetic field?  Diamagnetic – electrons are all ___________ so the substance is NOT attracted to a magnetic field.  Paramagnetic – _______________ electrons cause the substance to attracted to a magnetic field. Is carbon paramagnetic? recall:

Carbon ↑↓ ↑↓ ↑ ↑ __ 1s 2s 2p 2p 2p o _____! It has ___ unpaired electrons.