Atomic Structure Flashcards
Atomic Number (Z)
defines an element- no. protons
Mass Number (A)
no. protons + electrons
define isotope
different atoms of same element w/ same no. protons, diff no. neutrons
describe properties of isotopes of an element:
same chemical properties bc chemical properties depend on no. electrons but diff. physical properties bc depend on mass no.
relative atomic mass
average of all naturally occurring isotopes of an element relative to 1/12 of mass of carbon-12
what is mass spectrometer
measures proportions of each isotope present in a sample of an element in its vaporized state
greater deflection in a mass spectrometer=
higher charge and lower mass
how are electrons arranged in an atom
in energy levels
max no. electrons in an energy level
2n squared
describe the emission spectrum of H
a line spectrum w/ only certain frequencies/ wavelengths of light present- lines converge at a higher frequency/ energy
evidence for electrons being in energy levels
the emission spectrum because each line occurs when an excited electron falls back to its ground state bc unstable (when energy supplied to atom = excited from ground state - higher energy level), releasing E which corresponds to a particular wavelength that shows up as a line in the spectrum
A jump to n=1
ultraviolet spectrum bc highest E
a jump to n=2
visible spectrum
define convergence limit
when lines merge to form a continuum- beyond this point electron can have any E + free to leave atom
define orbital
region of space w/ high probability of finding an electron- represents discrete E level
Pauli Exclusion principle
max no. electrons in an orbital=2 + electrons must have opposite spins
Hunds Rule
orbitals must be filled singly first
ionisation energy
minimum amount of energy required to remove an electron from an atom of an element in its gaseous state
describe relationship between energy of a photon and frequency of electromagnetic radiation
E = hv
Energy (J)= frequency of light x Planck’s Constant = 6.63 x 10-34)
how to work out wavespeed
wavespeed = frequency x wavelength
why is second ionisation E always higher than the first?
once 1 electron removed = +ve ion which attracts electrons more strongly than neutral atoms therefore more E required to remove it
when 1 electron = removed, less e-e repulsion therefore electrons pulled in closer to nucleus= strongly attracted and difficult to move
how does ionisation E change across and period and why?
ionisation E increases along a period bc:
- nuclear charge increases
- electrons all removed from same energy level so no change in electron shielding
- attractive force on electrons increases from left to right
why does boron have a higher ionisation energy than beryllium?
electron removed from boron atom = in 2p sublevel which is at a higher energy than 2s so less energy required to remove it
why does nitrogen have a higher ionisation energy than oxygen?
nitrogen has 3 electrons in the 2p orbital whilst O has 4 in 2p orbital. Therfore there is pair of electrons in the 2p of oxygen. These electrons repel each other so electron = easier to remove
why does ionisation energy decrease down a group?
shielding of outer electrons from nucleus so attraction is weaker and are more easily removed
how to do electronic configurations for transition metals:
fill in 4s subshell first. for ions, remove 4s electrons first. copper and chromium only have 1 electron in 4s subshell bc atoms = most stable when subshell= full or half full
arrange subshells in order of increasing energy:
s, p,d,f
