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Topic 2. Atomic Structure

2.1 The Nuclear Atom

Proton # = Atomic #

Mass # = Proton # + Neutron #

Assuming neutral charge: Electron # = Proton # = Atomic #

Conclusions from Rutherford's Gold Foil Experiment:

  • There is a positively charged nucleus

  • Majority of the atom is empty







Isotopes have the same proton number but different neutron number

  • Different mass number/atomic mass

  • Same chemical properties but different physical properties

Ions have the same proton number but different electron number​

A mass spectrometer measures the abundance of different atoms/molecules in a sample

  1. Atomization: gas molecules separated

  2. Ionization: by the electron beam

  3. Acceleration: gas is filtered through negative electric plates → deflects anions created from Step 2

  4. Deflection: amount of deflection depends on the mass/charge ratio and velocity of the charged particles

    • lighter and higher charge → greater deflection

  5. Detection

2.2 Electron Configuration

Important formulas:

E = hv

c = vλ

E: energy in the form of photons (J)

v: frequency (Hz or 1/s)

λ: wavelength (m)

h: Planck's constant 6.63 × 10⁻³⁴ (J/s)

C: speed of light 3 × 10⁸ (m/s)

The shorter the wavelength, the higher the frequency, and the more energy it contains.

Emission spectra: when electrons are excited to a higher energy level, and then return to a lower energy level, they release a photon of a specific energy, as shown by a specific frequency of light.

  • Can be either continuous spectra or line spectra (ex. hydrogen)

Absorption spectra: shows how much light is absorbed at each wavelength of radiation (continuous)

Emission spectrum of hydrogen

  • Energy is not released continuously → line spectrum

  • This line spectrum converges at high frequencies, which means higher levels/shells get closer together

  • Electrons moving back to the lowest energy states and over the longest distances release the highest E (shortest λ)

At n=∞, the electron is considered to be no longer part of the atom

Higher energy level = higher freq = higher E = lower λ

Lyman series: when electrons fall to the ground state n=1

  • Emits UV light

  • High E, high freq, short λ

Balmer series: when electrons fall to n=2

  • Emits visible light

  • Use c=vλ and E=hv to find λ, then use the color wheel to determine what color light is emitted

    • Make sure to convert m to nm when using the IB color wheel​

Paschen series: when electrons fall to n=3

  • Emits infrared light

  • Low E, low freq, long λ

When comparing arrows in the SAME series, check the length of the arrow (AKA distance traveled by the electron)

  • ex. longer arrows (blue) have greater E than shorter arrows (red)

When comparing arrows from DIFFERENT series, see whether it's UV, visible, or infrared light to determine which E is greater

  • ex. UV light has greater E than visible and infrared

Electron configuration notation

Electrons go in shells or (principle) energy levels

The max number of electrons an energy level (n) can hold is 2n²

The energy levels contain sub-levels s, p, d, and/or f

Aufbau Principle: electrons enter the lowest energy orbital available first


Hund's Rule: every orbital in a subshell is singly occupied before any one is doubly occupied (to maximize stability and symmetry); all electrons in singly occupied orbitals spin in the same direction


Pauli Exclusion Principle: paired electrons must have opposite spins to reduce mutual expulsion

Exceptions to Aufbau:

  • ₂₄Chromium: 1s² 2s² 2p⁶ 3s² 3p⁶ 4s¹ 3d⁵

  • ₂₉Copper: 1s² 2s² 2p⁶ 3s² 3p⁶ 4s¹ 3d¹⁰

  • Because having one electron in each of the five d orbitals is more stable than having 4 filled and 1 empty/partially empty

Electron config. for ions: changes according to gain/loss of electrons

For transition metal ions: the electrons in 4s² are lost first

Noble gas configuration example: Titanium: [Ar] 4s¹ 3d²

Spherical shape

Dumbell shape

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