Chapter 2: Atomic Structure Flashcards
State and explain the 3 sub-atomic particles an atom is made of
ATOM: Positively-charged nucleus (nucleons: protons and neutrons) and negatively-charged electrons orbiting the nucleus
- Exception: Hydrogen (has 1 proton and 0 neutrons)
- Mass of atom: concentrated in the nucleus (relative mass of 1 for protons and neutrons vs relative mass of 5x10⁻⁴ for electrons)
- Volume of atom: concentrated in the electrons (size of nucleus vs size of atom)
Define atomic number (Z) and mass number (A)
ATOMIC NUMBER (Z): number of protons in an atom of that element
- Unique to each element and defines it
Atom: total charge is 0
- No. of protons = no. of electrons
MASS NUMBER (A): number of nucleons (protons + neutrons)
- No. of neutrons = A - Z
Define ions
ION: when an atom gains/loses electrons to become charged
- CATION: when atom loses electrons to become positively charged
- ANION: when atom gains electrons to become negatively charged
Define isotopes and discuss their properties
Mass number (A): Usually a decimal, as it is the average mass of the isotopes found in a naturally occurring sample of that element
ISOTOPES: atoms of the same element (same atomic/proton number) with a different mass number (different no. of neutrons)
- Same chemical properties: same no. of electrons (chemical properties determined by number and arrangement of electrons)
- Different physical properties: different no. of neutrons and different mass (physical properties determined by how fast an atom moves)
Calculate relative atomic mass (Aᵣ) and percentage abundance of isotopes
Aᵣ = [ (%abundance of Isotope A x mass of Isotope A) + (%abundance of Isotope B x mass of Isotope B)] / 100
State how isotopic composition is measured
Mass spectrometer: produces mass spectrums with a peak for each isotope proportional to the number of atoms of each isotope in the sample tested
Explain how electrons are arranged in an atom
Arranged in energy levels (shells with principal quantum numbers n, increasing in energy as they get further from the nucleus)
- Maximum number of electrons per energy level = 2n²
- Electrons fill shells from lowest to highest energy (closest to nucleus to furthest)
[Exception: Energy level 3 is only filled up to 8 electrons before energy level 4 starts to be filled]
Describe the electromagnetic spectrum
Light is represented by the electromagnetic spectrum
Frequency ∝ 1/Wavelength
Frequency ∝ Energy
- Increasing frequency/energy = Decreasing wavelength
Explain what an atomic emission spectrum is
Spectrometer: separates the various wavelengths of light emitted from a gas, generating an atomic emission spectrum (unique to every element)
Emission spectrum:
- Line spectrum: consists of sharp, bright lines on a dark background, where only some frequencies/wavelengths are present
- Continuous spectrum: all the colours merge together and all frequencies/wavelengths are present
Describe the atomic emission spectrum of hydrogen
Line spectrum: Red, Cyan, Indigo, Violet (converging at increasing frequency/decreasing wavelength)
Explain how emission spectra arise
When electric discharge is passed through a gas, the electron gains energy and is promoted to an energy level higher than its ground state (excited state). However, because the electron is unstable at the higher energy level, it falls back down to a lower energy level, where the extra energy gives out a photon of light that contributes to a line in the emission spectra.
- Eg. Hydrogen: Lines arise in the atomic emission spectrum of hydrogen when an electron in a hydrogen atom that gains energy and has been promoted to a higher energy level falls back to energy level __, producing a line in the ___ region of the emission spectrum.
Explain how the emission spectrum is proof of electrons existing in energy levels
A line spectrum being produced proves that electrons are in energy levels, as they can only transition between 2 discreet energy levels.
If electrons could have any energy and transition between all energy levels, a continuous spectrum would be formed instead
Explain how the emission spectrum of hydrogen is formed
Observed in the infrared, visible and ultraviolet regions:
- Infrared region (Paschen series) –> energy level 3
- Visible region (Balmer series) –> 2
- Ultraviolet region (Lyman series) –> 1
- Lines converge at higher frequencies in all regions
- At the convergence limit, lines merge to form a continuum (electron is no longer in the atom and can possess any energy)
Explain sub-energy levels and orbitals
Each main energy level (n) is made up of sub-energy levels:
- n=1 (max. no. of electrons: 2) : 1s²
- n=2 (max. no. of electrons: 8): 2s², 2p⁶
- n=3 (max. no. of electrons: 18): 3s², 3p⁶, 3d¹⁰
- n=4 (max. no. of electrons: 32): 4s², 4p⁶, 4d¹⁰, 4f¹⁴
Without any main energy levels, the energy levels of the sub-shells are:
s<p<d<f
- However, with the main energy levels, there may be reversals of orders between sub-levels
Determine the full electron configuration of atoms with up to 36 electrons (using Aufbau principle)
Eg. Fe: Period 4, Group 6
Aufbau principle: electrons fill sub-levels from lowest to highest energy
- 1s², 2s², 2p⁶, 3s², 3p⁶, 4s², 3d¹⁰,4p⁶, 4d¹⁰, 4f¹⁴
*4s² is filled before 3d¹⁰, as it is lower in energy
Alternative method: using periodic table
- Period number: highest main energy level filled
- Group number (-10): number of electrons in last filled sub-shell
- s-block (group 1-2)/d-block (group 3-12)/p-block (group 13-18): last filled sub-shell
Eg. Fe: Group 6 (d-block, 6 electrons in last filled shell), Period 4 (highest main energy level filled: 4)
- Last sub-shell filled: 3d⁶
