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7 .1: Electromagnetic Radiation PowerPoint Presentation

7 .1: Electromagnetic Radiation

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7 .1: Electromagnetic Radiation

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- Properties of Electromagnetic Waves:
- Wavelength ( λ ) :
- Distance Between two consecutive peaks or troughs in a wave
- Measured in meters

- Frequency (ν)
- Number of waves that pass a given point per second
- Measured in Hertz

- Speed (c)
- Speed of light
- Measured in meters/ second

- Relationship Between Properties
- Shortest wavelength = highest frequency
- Longest wavelength = lowest frequency
- INVERSE RELATIONSHIP

Types of Electromagnetic Radiation: Gamma, X-ray, UV, visible, IR, micro, radio

Einstein’s Photoelectric Effect

Max Plank & Quantum Theory: Energy is gained/lost in whole numbers multiples of the quantity hv ( frequency=v, Planck’s constant =h)

Planck’s Constant: h = 6.62606957 × 10-34 m2 kg / s ( J/s)

Planck discovered that energy is transferred to matter in packets of energy called quantum, rather than energy of matter being continuous.

ΔE = hν : The quantum of energy can be calculated from this equation

Phenomenon in which electrons are emitted from the surface of a metal when light strikes it

His observations are explained by assuming electromagnetic radiation is quantized (photons) and the threshold frequency is the minimum energy required to remove the electron.

- Dual Nature of Light:
- Light travels through space as a wave
- Light transmits energy as a particle
- Particles have wavelength, exhibited by diffraction patterns

- De Broglie’s Equation: Allows calculation of wavelength for a particle
- λ = h/mv
- Diffraction: results when light is scattered from a regular array of points or lines
- Diffraction Patterns: The interference pattern that results when a wave or a series of waves undergoes diffraction, as when passed through a diffraction grating or the lattices of a crystal. The pattern provides information about the frequency of the wave and the structure of the material causing the diffraction.

Continuous Spectrum: results when white light is passed through a prism. Contains all wavelengths of visible light

Line Spectrum: only see a few lines, each of which corresponds to discrete wavelength when passed thorough a prism. (Hydrogen emission spectrum)

Quantum Model : electron in a hydrogen atom moves around the nucleus only in certain allowed circular orbits

Bright line spectra confirms that only certain energies exist in the atom, and atom emits photons with definite wavelengths when the electron returns to a lower energy state.

Energy levels available to the electron in the hydrogen atom:

n= an integer

z= nuclear charge

J= energy in Joules

- Calculating the energy of the emitted photon
- Calculate electron energy in outer level
- Calculate electron energy in inner level
- Calculate the change in the energy
- ΔΕ= energy of final state- energy of initial state

- hc/ ΔΕ : to calculate the wavelength of emitted photon
- Energy Change in Hydrogen atoms
- Calculate the energy change between any two energy levels:

- Limitations of the Bohr Model
- Bohr’s model does not work for atoms other than hydrogen
- Electrons do not move in circular orbits

- Physical Meaning of ψ
- Square of the function is the probability of finding an electron near a particular point
- Represented as a probability distribution
- aka electron density map, electron density, electron probability

Electron bound to nucleus similar to standing waves

The exact path of the electron is not known

Heisenberg Uncertainty Principle- a limitation to the position and momentum of a particle at a given time

Radial Probability Distribution

Since the orbital size cannot be calculated, the size of the orbital is the radius of the sphere that an electron is in for 90% of the time

Principal quantum number (n)

- Main energy level
- 1, 2, 3, …
- Size and energy of orbital
- When n increases: orbital becomes larger, electron is further from the nucleus, higher energy b/c electron is less tightly bound to the nucleus so the energy is less negative

Angular momentum quantum number/Azimuthal QN (l)

- Sublevels, subshell
- 0...n-1 for each value of n
- Shape of atomic orbitals

Magnetic quantum number ()

- Integral values from l to –l
- Orientation of the orbital in space

s Orbitals

Spherical shape

Nodes for s orbitals of n=2 or greater

p Orbitals

Two lobes each

Occur in levels n=2 and greater

Each orbital lies along an axis

d Orbitals

Occur in levels n=3 or greater

Four orbitals with four lobes each centered in the plane indicated in the orbital label

Fifth orbital has two lobes along z axis and a belt centered in the xy plane

f Orbitals

Occur in levels n=4 and greater

Complex shapes

Usually not involved in bonding in compounds

Size of orbital:

Defined as the surface that contains 9-% of the total electron probability.

As n increases orbitals of the same shape grow larger.

Orbital Energies

- All orbitals with the same value of n have the same energy for hydrogen atoms (Degenerate)
- The lowest energy state = ground state
- When the atom absorbs energy the electrons can move to higher energy orbitals
- “excited state”

- Electron Spin Quantum Number
- An orbital can only hold two electrons, must have opposite spins.
- Spin can have +1/2 or -1/2

- Pauli Exclusion Principal
- In a given atom no two electrons can have the same set of four quantum numbers

Polyatomic Atoms: Atoms with more than one electron

3 energy contributions must be considered in description of the atom:

- Kinetic energy of electrons as they move around the nucleus
- Potential energy of attraction between nucleus and electrons
- The potential energy of repulsion between the two electrons

Electron correlation problem: Electron pathways are not known, so electron repulsive forces cannot be calculated exactly

- Average repulsions are approximated by...
- Treat each electron as it were moving in a field of charge that is the net result of the nuclear attraction and average repulsions of all other electrons

Screening or Shielding

Electrons are attracted to the nucleus

Electrons are repulsed by other electrons

Electrons would be bound more tightly if other electrons weren’t present

Closer proximity to the nucleus = lower energy

Originally constructed to represent patterns observed in chemical properties of elements

Mendeleev and Meyer both independently conceived present periodic table

Mendeleev also corrected several atomic masses

- Periodic Table Vocab:
- Valence electrons: electrons in outermost principal quantum level of an atom
- Transition metals: “d” Block
- Lanthanide and Actinide Series : “f” block
- Representative Elements: Group 1A through 8A
- Metalloids: Border between metals and nonmetals, exhibit properties of both

Aufbau Principle: “As protons are added one by one to the nucleus to build up elements, electrons are similarly added to these hydrogen like orbitals”

Hunds Rule: “The lowest energy configuration for an atom is the one having the maximum number of unpaired electrons allowed by Pauli principle in a particular set of degenerate orbitals

7.13: The Properties of a Group: Alkali Metals

Ionization energy: energy required to remove an electron from an atomic (increase across period, decreases with increasing atomic number within a group)

Electron affinity: energy change associated with the addition of an electron (decrease down period, increase across period)

Atomic Radius: Determination of radius (increases down group, decreases across period)

- Easily lose valence electrons
- Reducing agents
- React with water
- Large hydration energy
- Positive ionic charge