 Arrangement of Electrons in Atoms

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Arrangement of Electrons in Atoms
Chapter 4

The New Atomic Model Investigations  relationship between light and atom’s electrons How are electrons arranged? Why don’t they fall into the nucleus?

Light a wave or particle?
Wave Description: Electromagnetic Radiation: energy that acts like a wave in space All forms create Electromagnetic Spectrum

Electromagnetic Spectrum

Electromagnetic Spectrum
All forms move at speed of light, c, 3.00x108 m/s Forms identified by: wavelength, , the distance b/ corresponding points on adjacent waves. Units: nm, cm, or m Frequency, , # of waves that pass a given point in a specific time, 1 sec. Unit: 1/s = Hertz, Hz

Wavelength and Frequency

Wavelength and Frequency
Inverse proportion equation!! Frequency, 1/s speed of light, m/s wavelength, m

Calculation Calculate the wavelength of a radio wave with a frequency of x 106s-1 Determine the frequency of light whose wavelength is nm.

Particle Nature of Light
Photoelectric Effect: emission of electrons from a metal when light shines on the metal

Photoelectric Effect Light had to be certain frequency to knock e- loose Wave theory  any frequency should work (just might take a while) Light must also be a particle! Max Planck(1900) explanation: objects emit energy in small packets called quanta Video - 16

Max Planck Quantum of energy is the smallest amount of energy that can be lost or gained by an atom E = h Frequency, s-1 Energy of quantum, in joules, J Planck’s constant, 6.626x10-34 Js

Energy Calculation What is the energy of green light, with a wavelength of 500. nm?

Albert Einstein Light is both wave and particle!
Particle of light = photon, having zero mass and a quantum of energy Photons hit metal and knock e- out, but photon has to have enough energy

H-atom Emission Spectrum
Pass a current through gas at low pressure it excites the atoms Ground state: lowest energy state of an atom Excited state: atom has higher potential energy than it has in ground state

H – Atom Spectrum When atom jumps from excited state to ground state it gives off energy  LIGHT! E2 Ephoton = E2 – E1 = hv E1

Bohr Model of H-atom

H-atom Line Emission Spectrum

Element Emission Spectras
Helium – 23 lines Neon – 75 lines Argon lines Xenon – 139 lines Mercury – 40 lines

H-atom Line Emission Spectrum
More lines in UV (Lyman series) and IR(Paschen series) Why did H-atom only emit certain colors of light? Explanation led to new atomic theory  Quantum Theory

Bohr Model of H-atom 1913 – Niels Bohr
e- circles nucleus in certain paths, orbits or atomic energy levels e- is higher in energy the farther away from nucleus e- cannot be between orbits Video - 23

Bohr Model of H-atom

Bohr Model of H-atom From wavelengths of emission spectrum Bohr calculated energy levels of H-atom Model worked ONLY for H-atom

Quantum Model of Atom Can e- behave as a wave?
Yes! To find e- use a photon, but photon will knock the e- off course Heisenberg Uncertainty Principle: impossible to determine position and velocity of a particle at the same time.

Schrödinger Wave Equation
1926 – developed equation and only e- waves of certain frequencies were solutions Quantization of e-  probability of finding e- in atom No neat orbits  probability clouds or orbitals

Electron Configurations

Atomic Orbitals Def: 3-D region around nucleus that indicates the probable location of an electron Energy levels or shells: Numbered 1-7 Smaller number = closer to nucleus, lower energy

Sublevels Each shell has sublevels s p d f 1 – s orbital
3 – p orbitals d 5 – d orbitals f 7 – f orbitals

Shells and Sublevels Shells and sublevels together: 1s 2s, 2p
3s, 3p, 3d 4s, 4p, 4d, 4f, etc. s is the lowest energy and f is the highest

Orbitals Each orbital in a sublevel can hold a maximum of 2 e-
1 – s 2 e- max. 3 – p orbitals 6 e- max. 5 – d orbitals 10 e- max. 7 – f orbitals 14 e- max.

Electron Configurations
Arrangement of e- in atom Orbital Notation: H has 1e- Rules: Aufbau Principle: electron occupies lowest energy level that can receive it

Electron Configurations
2. Pauli Exclusion Principle: no two e- in an sublevel orbital can have the same spin 3. Hund’s Rule: orbitals of equal energy are occupied by one e- before pairing up e-. All single occupied orbitals must have same spin. He – 2e-

Energy of sublevels

Electron Configurations

Electron Configuration Notation
B Ni Hg

Noble Gas Notation Use noble gas from previous row Al Pb

Special Cases d sublevel more stable with half-filled or completely filled sublevel Cr Cu

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