Atomic Structure Flashcards
John Dalton
Thought atoms were spherical.
Different spheres constituted different elements.
Thought atoms were fundamental units of matter and were indivisible.
J.J. Thompson
Discovered atoms were divisible.
Discovered and measured mass of electron.
Led to Plum Pudding Model (atom was a positively charged sphere with negatively charged electrons embedded in it)
Ernest Rutherford
Fired alpha particles at a thin gold film.
Most particles passed straight through so disproved Plum Pudding Model.
Led to Nuclear Model (atom was mostly empty space with positive nucleus and orbiting electrons)
Neils Bohr
Developed first quantum theory and discovered electron shells.
When electrons move between orbits, electromagnetic radiation of particular frequency must be emitted or absorbed
Atomic Mass Unit (amu)
1/12 of mass of single carbon-12 atom
Masses of subatomic particles
Proton = 1.0073 amu Neutron = 1.0087 amu Electron = 0.00055 amu
Isotopes of hydrogen
Protium: one proton, one electron, no neutrons
Deuterium: one proton, one electron, one neutron
Tritium: one proton, one electron, two neutrons
Relative isotopic mass
Average mass of a single isotope of an element divided by 1/12 of the mass of an atom of carbon-12
Relative atomic mass
Average mass of a single atom of an element divided by 1/12 of the mass of an atom of carbon-12
Relative molecular mass
Average mass of a molecule divided by 1/12 of the mass of an atom of carbon-12
Time of flight mass spectrometer
Ionisation Acceleration Ion drift Ion detection Data analysis
Ionisation
electrospray/ electron impact
Acceleration
Electric field applied to give all ions same kinetic energy.
Heavier particles move slower.
Ion drift
Ions deflected by magnetic field into a curved path.
Radius of their path is dependant on the charge and mass of the ion.
Ion detection
Positive ions hit detection plate, gain electron producing flow of charge.
Greater abundance produces greater current.
Data analysis
Analysed and recorded as mass spectra.
Electrospray
Sample is dissolved in volatile, polar solvent
Injected through fine hypodermic needle giving fine mist or aerosol
Tip of needle has high voltage
At tip of needle sample molecule, M, gains a proton, H+ from solvent forming MH+
Solvent evaporates away while the MH+ ions move towards a negative plate
Peak with largest m/z value is equal to mass of MH+, have to subtract 1 to get Mr of molecule
Used for larger organic molecules, fragmentation does not occur
Electron Impact
Vaporised sample is injected at low pressure
Electron gun fires high energy electrons at sample
Knocks out an outer electron
Forms positive ions with different charges
Used for molecules with low formula mass as fragmentation may occur with larger molecules
Mass spectra for Cl2 and br2
Cl has two isotopes: Cl35 (75%) Cl37 (25%) Br has two isotopes Br79 (50%) Br81 (50%)
Use of mass spectrometers
Included in planetary space probes so elements on other planets can be identified as elements may have different composition of isotopes
Electron configuration
Energy levels get closer together further from the nucleus.
Each orbital holds up to two electrons. Spin in opposite directions to minimise repulsion.
4s orbital fills before 3d. When d block elements form ions, lose 4s electrons first.
degenerate
Orbitals with the exact same energy
isoelectronic
Particles with same electron configuration
electron configuration of copper
Cu = 1s22s22p63s23p64s13d10
electron configuration of chromium
Cr = 1s22s22p63s23p64s13d5
Simple ions
Name of positive ion is same as name of atom
Na = sodium
Na+ = sodium ion
Name of negative ion is atom stem with -ide on the end
Br = bromine
Br- = bromide ion
Shapes of orbitals
S subshells are spherical
P subshells are shaped like dumbbells
First ionisation energy
Energy required to remove one mole of electrons from one mole of gaseous atoms to form one mole of gaseous 1+ ions
Second ionisation energy
Energy required to remove one mole of electrons from one mole of gaseous ions with a single positive charge to form one mole of gaseous 2+ ions
Factors affecting ionisation energies
Atomic radius
Nuclear charge
Shielding
Distance
Atomic radius effect on ionisation energy
higher the atomic radius
lower the ionisation energy
Nuclear charge effect on ionisation energy
greater number of protons
greater nuclear charge
higher ionisation energy
Shielding effect on ionisation energy
more inner electrons causes
more shielding
lower ionisation energy
Distance effect on ionisation energy
Electrons further away from nucleus
easier to remove
lower ionisation energy
Trends down a group
ionisation energy
Outer electron further from nucleus, higher atomic radius Increased shielding due to more inner electrons Nuclear charge increases Ionisation energy decreases
Trends across a period
ionisation energy
Electrons being added to same shell, atomic radius decreases Higher nuclear charge Same shielding Ionisation energy increases
Group 3 first ionisation energies
Lower than expected first ionisation energy Have s2p1 arrangement outer p1 electron further from nucleus Inner s2 electrons increase shielding Less energy required
Group 6 first ionisation energies
Lower than expected
Have a p4 arrangement
Repulsion of two electrons in same p orbital
Less energy required
Group 1 first ionisation energies
Lowest first ionisation energy in every period
Greatest atomic radius
Lowest nuclear charge
Group 0 first ionisation energies
Highest first ionisation energy in every period
Smallest atomic radius
Highest nuclear charge
Patterns in second ionisation energy
Patterns in first ionisation energy shifted one to the left
Group 1 elements would have highest second ionisation energy
