AP Chemistry

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7. Atomic structure and periodicity

Electromagnetic radiation

electromagnetic radiation: waves of energy that have magnetic

electricity + magnets + releasing

e.g. microwave, satellite, visible light (rainbow), ultraviolet, infrared, radio, television, gamma rays, X rays

wavelength (λ): distance from one peak to another peak (meters/m)

frequency (ν): the number of waves that pass a location in one second (hertz/Hz/s⁻¹)

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visible light: 400 nm - 700 nm

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c: speed of light (2.9979 × 10⁸ m/s)

Nature of matter

Max Planck:

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E: energy (J)

h: Planck’s constant

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ν: frequency (s⁻¹)

energy is quantized (little packets)

Albert Einstein: photoelectric effect (energy shot at metal → electrons fly out)

only certain wavelengths make electrons come off

visible/ultraviolet light: electrons change energy levels (think Bohr model)

microwaves: molecule rotates (think microwave plate)

infrared: molecules vibrate

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E: energy

m: mass

c: speed of light

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dual nature of matter: matter can be both a particle and a wave

de Broglie: all matter are waves

Atomic spectrum

if you put atoms in a prism, only certain colors appear

light spectrum

Bohr’s model of the atom

⁠

if you add energy to an orbital:

the electron will jump up an orbital/energy level

electron returns to original orbital

energy released (often visible light)

works for hydrogen (1 electron), but not with other elements

energy levels (n): orbitals according to Bohr model

n = 1: ground state

Quantum mechanical model

Ervin Schrödinger: equation for electron probability

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Werner Heisenberg: we can know where an electron is or the momentum of the electron, never both

Quantum numbers and orbitals

quantum numbers: “address” of electron (big → small)

principal quantum number (n): 1 → ∞

energy level

angular momentum number (ℓ): 0 → n-1

0 = s

1 = p

2 = d

3 = f

4 = g

magnetic quantum number (m_ℓ): -ℓ → ℓ

1 orbital can hold two electrons

s: 2 electrons

p: 6 electrons

d: 10 electrons

f: 14 electrons

electronic spin (mₛ): 1/2 or -1/2

electron symbol: ⥮

Wolfgang Pauli: Pauli exclusion principle

every electron in an atom can only have one address

cannot have two electrons with the same address

Polyelectronic atoms

aufbau principle (“build up”): electron will always go to the lowest energy level first

examples

1 electron → 1 energy level

1 shape → s (orbital)

1 electron

1s¹ (electron configuration)

2 electrons → 1 energy level

1 shape → s

2 electrons

1s²

3 electrons

1s²2s¹

4 electrons

1s²2s²

5 electrons

1s²2s²2p¹

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periodic table

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look at row number

check orbital group

Periodic trends

atomic radius: distance from the nucleus to the valence electrons in an atom

down: increases (more shells)

right: decreases (more protons → more attraction)

ionization energy: amount of energy needed to remove one electron

down: decreases (electrons further from nucleus)

right: increases (want to lose on left but not on right)

electron affinity: amount of energy that is released when an electron is added

down: depends on column but generally decreases (electrons further away)

right: increases release of energy (more stable)

ionic radius: distance from the nucleus to the valence electrons in an ion

down: increases (more shells)

right

left: decreases (losing electrons)

goes up (more electrons)

right: decreases (adding electrons)

Electromagnetic radiation

Nature of matter

Atomic spectrum

Bohr’s model of the atom

Quantum mechanical model

Quantum numbers and orbitals

Polyelectronic atoms

Periodic trends

Gallery

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