Chemistry - Chapter 6-7: Electron Configuration, EM Radiation

Atomic structure and periodic trends (more focused on vocab words. Look here for further detailing on periodic trends: https://www.notion.so/mahalkita/Section-7-Periodic-Trends-1cd5fbabd389424cb77ef34bed624812)

c=vλ

c=vλ

represents the relationship between wavelength, frequency, and the speed of light

c = speed of light (2.998×10⁸ m/s)

v = frequency (hertz or sec^-1)

λ = wavelength (meters)

Wavelength and frequency are inversely related

6.626×10^-34 J/s

Planck’s constant.

Represented by h

EM spectrum

contains, in the following order with increasing energy,

• radio

• microwave

• infrared

• visible light

• ultraviolet

• x-ray

• gamma ray

Einstein’s photoelectric effect

electrons are emitted from the surface of a metal from light contact → must first reach minimum critical frequency of light (v_o) to be ejected

amount of electrons ejected depends on intensity of light.

energy of electrons will increase if there is an increased frequency

Electron Absorption Motion/Transitions

transitions that are perceived with the absorption of energy from electrons

Rotational Spectrum

absorption of microwave and far infrared radiation

• rotates from heat

Vibrational spectrum

absorption of infrared and high microwave radiation

• symmetric/asymmetric stretching and bending of bonds

Electronic spectrum

absorption of visible light and ultraviolet radiation

• wavelength shortens the bigger the jump

• the falling back of excited e- produces visible light

Wave-Particle Duality

light moves as a wave, but is comprised of particles

λ=h/mv

De Broigle’s equation that allows calculation of wavelength of any object

• Based off his ideas that all matters have wave-like properties,

  • the larger the object (m = mass in kg), the more miniscule the wavelength, and therefore, its visible effect on the object

Bohr’s Model

revolutionary in the idea that electrons only exist in specific energy states (we can only define the probability of finding an electron at a given location)

Heisenberg’s Uncertainty Principle

The more certain you are about where electron is, the less certain you can be about its energy, and vice versa

Principal Quantum Number

energy level/shell. Represented by n, in which:

• n² = orbitals in energy level

• 2n² = electrons in energy level

Angular Momentum Quantum Number

sublevel and or subshells. Describes the shape of an orbital. Either: s, p, d, f

Magnetic Quantum Number

orientation of electrons in orbitals

Paramagnetism

unpaired electrons, responds to magnetic field

Ferromagnetic

the spins of unpaired electrons aligned in the same direction, even in the absence of a magnetic field

Penetration Effect

the probability of electrons spending their time near the nucleus, in relation to their distance

• explains why certain orbitals fill first → closer to nucleus

• further from nucleus → more likely to lose electrons

effective nuclear charge (Z_eff)

effect of the nucleus on valence electrons

• result of nuclear charge, offset by:

  • electron shielding

  • distance from nucleus

Coulomb’s Law

force of attraction between nucleus and electrons decrease exponentially with distance

Ionization Energy

energy required to knock off an atom. Equal to E_photon - KE_electron, measured via PES (values are a known energy of photons that strike a substance to produce a measured KE)

Photo-Electro Spectroscopy

PES. Shows relative intensities of electrons (peaks) that correspond to the substance’s electron configuration [y-axis], over ionization energy required of each electron cluster in relation to their distance to nucleus [x-axis].

• shows valence e- takes less energy to eject than those closer to nucleus

electron affinity

energy change associated with the addition of an electron

electronegativity

measure of the ability of an atom to attract shared electrons