Chapter 10- Part One Modern Atomic Theory Objectives: Review history… (10.1) Describe electromagnetic radiation (10.2) Describe the Bohr atom (10.3) Explain energy levels of electrons and diagram atomic structures for elements (10.4 & 10.5)

Review… Dalton Thomson Rutherford Model doesn’t explain how the negative electron can stay in orbit and not be attracted to the positive proton

Electromagnetic Radiation Light travels in waves, similar to waves caused by a moving boat or a pebble tossed in a pond Light is a form of Electromagnetic Radiation Form of energy that exhibits wavelike behavior as it travels through space

Electromagnetic Radiation All waves can be described in 3 ways: Amplitude – the height of the wave, results in the brightness or intensity of the light Wavelength (l): distance between consecutive peaks in a wave Frequency (n): number of waves that pass a given point in a second

Electromagnetic Radiation Speed of light in air: Electromagnetic radiation moves through a vacuum at speed of 3.00 x 108 m/s Since light moves at constant speed there is a relationship between wavelength and frequency: c = ln Wavelength and frequency are inversely proportional C = speed of light L = wavelength V = frequency

Electromagnetic Spectrum Visible spectrum – white light (what we can see) Wavelengths of visible spectrum are 350 nm to 700 nm

Quantum Theory Wave theory does not explain Heated iron gives off heat 1st red glow yellow glow white glow How elements such as barium and strontium give rise to green and red colors when heated

Energy is released in Quanta Quantum Theory Max Planck (1858-1947) Proposed that there is a fundamental restriction on the amounts of energy that an object emits or absorbs, and he called each of these pieces of energy a quantum. Energy is released in Quanta

E = hn E = energy Quantum Theory A quantum is a finite quantity of energy that can be gained or lost by an atom E = hn E = energy v = frequency h = 6.626 x 10-34 J/s This constant, h, is the same for all electromagnetic radiation

Photoelectric Effect The emission of electrons by certain metals when light shines on them Albert Einstein (1905) used Planck’s equation to explain this phenomenon; proposed that light consists of quanta of energy that behave like tiny particles of light Photon = individual quantum light (also known as a particle of radiation)

Photoelectric Effect He (Einstein) explained that the photoelectric effect would not occur if the frequency and therefore the energy of each photon is too low to dislodge an electron. Analogy: 70 cents placed in soda machine: no soda 30 cents more and you will get your soda

Now… Light can be described as both particles and waves Dual Wave-Particle Nature of Light was accepted What does this mean for the atom???

Line Spectrum Elements in gaseous states give off colored light High temperature or high voltage Always the same Each element is unique http://home.achilles.net/~jtalbot/data/elements/ Also known as the “atomic emission spectrum of an element” Visible light or electromagnetic radiation is emitted when an atom passes from a state of higher potential energy to a state of lower potential energy

Line Spectrum Ground state Excited state Lowest energy level available State in which electron has a higher potential energy than in its ground state Farther from nucleus Higher potential energy

Line Spectrum Electron falls from higher energy level to lower one…emits light at a specific frequency Color of light emitted depends on difference between excited state and ground state See figure 10.5 page 201

Line Spectrum Each band of color is produced by light of a different wavelength Each particular wavelength has a definite frequency and has definite energy Each line must therefore be produced by emission of photons with certain energies

Line Spectrum

Line Spectrum Whenever an excited electron drops from such a specific excited state to its ground state (or lower excited state) it emits a photon The energy of this photon is equal to the difference in energy between the initial state and the final state.

Niels Henrik David Bohr 1885-1962 Physicist Worked with Rutherford 1912 Studying line spectra of hydrogen

Niels Henrik David Bohr 1913 – proposed new atomic structure Electrons exist in specific regions away from the nucleus Electrons revolve around nucleus like planets around the sun

The Bohr Atom Nucleus with protons and neutrons Electrons move in “stationary states” which are stable (paths or orbits) When an electron moves from one state to another the energy lost or gained is done is ONLY very specific amounts Each line in a spectrum is produced when an electron moves from one stationary state to another

The Bohr Atom Model didn’t seem to work with atoms with more than one electron Did not explain chemical behavior of the atoms Safeco field example

Wave Matters… Louis de Broglie (1924) Proposed that electrons might have a wave-particle nature Used observations of normal wave activity Safeco field example

Wave Matters… Erwin Schrodinger (1926) Used mathematical understanding of wave behavior – devised an equation that treated electrons moving around nuclei as waves Quantum Theory Safeco field example

Quantum Theory Describes mathematically the wave properties of electrons and other very small particles Applies to all elements (not just H)

Energy Levels of Electrons Principal energy levels Designated by letter n Each level divided into sublevels 1st energy level has 1 sublevel 2nd energy level has 2 sublevels Etc.

Energy Levels of Electrons

Orbitals Electrons don’t actually orbit like planets Orbital: region in space where there is a high probability of finding a given electron Each orbital sublevel can hold 2 electrons Safeco field example

Orbitals Each sublevel (orbital) has a specific shape Safeco field example http://daugerresearch.com/orbitals/

Orbitals Pauli exclusion principle: an atomic orbital can hold a maximum of two electrons which must have opposite spins Electrons can only spin in two directions Shown with arrows Safeco field example

Rules for Orbital Filling Pauli’s Exclusion Rule No two electrons have the same set of quantum numbers Hund’s Rule Electrons will remain unpaired in a given orbital until all orbitals of the same sublevel have at least one electron 1s 2s 2p 3s 3p

Rules for Orbital Filling Diagonal Rule The order of filling once the d & f sublevels are being filled Due to energy levels

Rules for Orbital Filling Safeco field example

Quantum Numbers Numbers that specify the properties of atomic orbitals and their electrons Principle Quantum Numbers: Symbolized by n, indicates the main energy levels surrounding a nucleus, which indicates the distance from the nucleus (shells or levels)

Quantum Numbers Orbital Quantum Number: Indicates the shape of an orbital (subshell or sublevels) s, p, d, f Principal Quantum # Orbital Quantum # 1 1s 2 2s, 2p 3 3s, 3p, 3d 4 4s, 4p, 4d, 4f

Quantum Numbers Magnetic Quantum Number: Indicates the orientation of an orbital about the nucleus Orbital position with respect to the 3-dimensional x, y, and z axes

Quantum Numbers Spin Quantum Number: Indicates two possible states of an electron in an orbital Type of Orbital Number of Orbitals s 1 ( ) p 3 (x, y, z) ( , , ,) d 5 ( , , , , ) f 7 Each orbital holds a maximum of 2 electrons

Application of Quantum Numbers Several ways of writing the address or location of an electron Lowest energy levels are filled first Electron Configuration: using the diagonal rule, the principal quantum number (n), and the sublevel write out the location of all electrons 12C: 32S: 1s22s22p2 1s22s22p63s23p4

Application of Quantum Numbers Orbital filling electron diagram: using Hund’s rule and the diagonal rule write out the location of all electrons See examples on whiteboard

Chapter 10 – Part Two The Periodic Table Objectives: Understand the arrangement of the Periodic Table (10.6) Identify connections between electron configuration and placement on the periodic table

The Periodic Table 1869 – arrangement proposed by Dmitri Mendeleev And Lothar Meyer (different layout) Still similar today Based on increasing atomic masses and other characteristics Was able to predict properties of elements not yet discovered….and was correct!

The Periodic Table Horizontal rows Vertical Columns Periods Corresponds to outermost energy level Vertical Columns Groups or families Similar properties; reactions

The Periodic Table Several systems for naming groups Left to right, 1-18 Roman numerals and A and B Used in this book Group A: Representative Elements Noble Gases IA – Alkali Metals IIA – Alkaline Earth Metals VIIA - Halogens Group B: Transition Elements

The Periodic Table Chemical behavior and properties of elements in a particular family similar Have the same outer shell electron configuration Figure 10.15 page 211 Noble gas configuration (shortcut) Use previous noble gas in square brackets Finish with valence electrons

The Periodic Table Examples: K is 1s22s22p63s23p64s1 or [Ar]4s1 Ca is 1s22s22p63s23p64s2 or [Ar]4s2 Write abbreviated configuration for the following elements: Fr Y

The Periodic Table Arrangement of Periodic Table also means that elements filling similar orbitals are grouped s block p block d block f block Know these blocks…

The Periodic Table - Highlights The number of the period corresponds to the highest energy level occupied by electrons in that period The group numbers for the representative elements are equal to the total number of valence electrons in that group

The Periodic Table - Highlights The elements of a family have the same outermost electron configuration (just different energy levels) The elements within each of the s, p, d, and f blocks are filling the corresponding orbitals There are some discrepancies with order of filling (not covered in this book)