Electrons in Atoms

 Wavelength ( ) - length of one complete wave measured in m, cm, or nm  In light it tells us which color it is  Frequency ( ) - # of waves that pass a point during a certain time period,  hertz (Hz) = 1/s  Amplitude (A) - distance from the origin to the trough (dip) or crest (peak  Amount of energy the wave is carrying - height of wave. It is measured in meters.  In SOUND it tells us how LOUD it is. In LIGHT it tells how BRIGHT it is.

A greater amplitude greater frequency crest origin trough A

 Understanding electronic structure of atoms  Must understand light  Emitted and absorbed by substances.  Visible light - type of Electromagnetic Radiation (EM)  Carries (radiant) energy through space  Travels at speed of light  Exhibits wavelike behavior.  Think of light as particle  help understand how EM radiation and atoms interact

LOWENERGYLOWENERGY HIGHENERGYHIGHENERGY

 Move through a vacuum at the ‘speed of light’ 3.00 x 10 8 m/s  Behaves like waves that move through water  Result of energy transferred to the water (from a stone)  Expressed as up and down movement  Both electric and magnetic properties

 Wave Speed = (distance between wave peaks) x (frequency) = (wavelength) x (frequency)  EM radiation moves through a vacuum at the “speed of light” 3.00 x 10 8 m/s also called c.  A lower energy wave (infrared and red) has a longer wavelength( ) and lower frequency(f)  A higher energy wave (blue - violet) has a shorter wavelength( ) and higher frequency(f).

 Frequency & wavelength are inversely proportional c = c:speed of light (3.00  10 8 m/s) :wavelength (m, nm, etc.) :frequency (Hz)

 EX: Find the frequency of a photon with a wavelength of 434 nm. GIVEN: = ? = 434 nm = 4.34  m c = 3.00  10 8 m/s WORK : = c = 3.00  10 8 m/s 4.34  m = 6.91  Hz

 Planck (1900)  Observed - emission of light from hot objects  Concluded - energy is emitted (absorbed or released) in small, specific amounts (quanta)  Quantum - smallest energy packet that can be emitted or absorbed as EM radiation by an atom.

E:energy (J, joules) h:Planck’s constant (  J·s) :frequency (Hz) E = h zPlanck proposed that the energy, E, of a single quantum energy packet equals a constant (h) times its frequency zThe energy of a photon is proportional to its frequency.

 EX: Find the energy of a red photon with a frequency of 4.57  Hz. GIVEN: E = ? = 4.57  Hz h =  J·s WORK : E = h E = (  J·s ) ( 4.57  Hz ) E = 3.03  J

 Energy is always emitted or absorbed in whole number multiples of hv, such as hv, 2 hv, 3 hv, 4hv, ….  The allowed energies are quantized  values are restricted to certain quantities.  The notion of quantized rather than continuous energies is strange.  Ramp vs Staircase  Ramp - vary the length your steps and energy used on the walk up.  Stairs - must exert exactly the specific amount of energy needed to reach the next step.  Your steps on steps are quantized, you cannot step between them.

 Planck (1900) vs. Classical TheoryQuantum Theory

 Einstein (1905)  Observed – photoelectric effect  Dispersed light falls on metal samples, the different frequencies produce different energetic photoelectrons

 Einstein (1905)  Concluded - light has properties of both waves and particles (photons) “wave-particle duality”  Photon - particle of light that carries a quantum of energy  Used planck’s quantum theory to deduced that: E photon = hv

Electrons in Atoms

ground state excited state ENERGY IN PHOTON OUT yElements’ atoms absorb electrical energy ye- get excited, become unstable, and release energy yEnergy is in form of light ySet of frequencies of EM waves

 e - exist only in orbits with specific amounts of energy called energy levels  Therefore…  e - can only gain or lose certain amounts of energy  only certain photons are produced  Ground state = lowest allowable atomic electron energy state  Excited state = any higher energy state

 Energy of photon depends on the difference in energy levels  Bohr’s calculated energies matched the IR, visible, and UV lines for the H atom

 Each element has a unique bright-line emission spectrum.  “Atomic Fingerprint” Helium  Examples:  Iron  Now, we can calculate for all elements and their electrons – next section

 Louis de Broglie (1924)  Proposed eˉ in their orbits behave like a wave EVIDENCE: DIFFRACTION PATTERNS ELECTRONS VISIBLE LIGHT

 Heisenberg Uncertainty Principle  Impossible to know velocity and position of an electron at the same time  Trying to observe an electron’s position changes its momentum  Trying to observe an electron’s momentum changes its position  Electrons cannot be locked into well-defined circular orbits around the nucleus.

 Orbital (“electron cloud”)  a specific distribution of electron density in space.  Each orbital has a characteristic energy and shape. Orbital

Specify the “address” of each electron in an atom UPPER LEVEL

1. Principal Quantum Number ( n = 1, 2, 3, …) (see periodic table left column)  Indicates the relative size and energy of atomic orbitals  As (n) increases, the orbital becomes larger, the electron spends more time farther from the nucleus  Each major energy level is called a principle energy level Ex: lowest level = 1 ground state, highest level = 7 excited state

2. Energy Sublevel  Defines the shape of the orbital (s, p, d, f)  # of orbital related to each sublevel is always an odd # s = 1, p = 3, d = 5, f = 7  Each orbital can contain at most 2 electrons s p d f

Subscripts x, y, z designates orientation  Specifies the exact orbital within each sublevel

pxpx pypy pzpz

4. Spin Quantum Number ( m s )  Electron spin  +½ or -½  An orbital can hold 2 electrons that spin in opposite directions.

 Pauli Exclusion Principle  A maximum of 2 electrons can occupy a single atomic orbital  Only if they have opposite spins 1. Principal #  2. Energy sublevel  3. Orientation  4. Spin #  energy level (s,p,d,f) x, y, z exact electron

Electron Configuration Electrons in Atoms

A. General Rules zAufbau Principle yElectrons fill the lowest energy orbitals first. y“Lazy Tenant Rule”

RIGHT WRONG A. General Rules zHund’s Rule yWithin a sublevel, place one e - per orbital before pairing them. y“Empty Bus Seat Rule”

O 8e - zOrbital Diagram zElectron Configuration 1s 2 2s 2 2p 4 B. Notation 1s 2s 2p

zShorthand Configuration S 16e - Valence Electrons Core Electrons S16e - [Ne] 3s 2 3p 4 1s 2 2s 2 2p 6 3s 2 3p 4 B. Notation zLonghand Configuration yValence electrons: determine chemical properties of that element & are the electrons in the atoms outermost orbital

© 1998 by Harcourt Brace & Company s p d (n-1) f(n-2) Notation

zShorthand Configuration yCore e - : Go up one row and over to the Noble Gas. yValence e - : On the next row, fill in the # of e - in each sublevel. Shorthand Notation

[Ar]4s 2 3d 10 4p 2 C. Periodic Patterns zExample - Germanium

zFull energy level zFull sublevel (s, p, d, f) zHalf-full sublevel D. Stability

zElectron Configuration Exceptions yCopper EXPECT :[Ar] 4s 2 3d 9 ACTUALLY :[Ar] 4s 1 3d 10 yCopper gains stability with a full d-sublevel. D. Stability

zElectron Configuration Exceptions yChromium EXPECT :[Ar] 4s 2 3d 4 ACTUALLY :[Ar] 4s 1 3d 5 yChromium gains stability with a half-full d-sublevel. D. Stability

zIon Formation yAtoms gain or lose electrons to become more stable. yIsoelectronic with the Noble Gases.

O 2- 10e - [He] 2s 2 2p 6 D. Stability zIon Electron Configuration yWrite the e - config for the closest Noble Gas yEX: Oxygen ion  O 2-  Ne