Structure 1.3.1, 1.3.2- Emission Spectra Flashcards
(30 cards)
Planck’s equation
E = hf
E = h (C/ wavelenght)
wavelnght = nm, convert to m
define ionization energy
the energy required for an atom to lose an electron and become a cation (1st IE)
->essentially the desire to become a cation, (a lower IE = more readily becomes a cation)
Define wavelength
the distance b/w two adjacent peaks in wave (reported in nm, to be converted to meters for calculations)
1nm = 1x10^-9 m
define frequency
of waves that pass a given point per second (s^-1 is preferred, but can be reported as Hz)
Colors in the visible spectrum (from lowest to highest energy)
ROY G BIV
What determines the color emitted on the visible spectrum
by the wavelenght that a photon emits, when the wavelenght is long we see reds, when short we see violets
what’s a photon
packet of energy
define continuous spectrum
-a spectrum that occurs when white light is passed through a prism
-shows all wavelengths or freqencies of visible light with a smooth transition b/w colors (rainbow)
define line spectrum
only has certain wavelengths or frequencies of light, produced by electrons in the excited state
how are hydrogen emission spectrum created
produced as electrons transition from a higher energy level to n=2-
how are hydrogen absorption spectrum created
created when a gaseous sample is electrified
How are line spectrums produced
They’ve are produced by electrons inside atoms and ions as they absorb and emit energy, and jump from a lower energy level to a higher energy level, and then return to their original position
What happens when an absorption + emission Spectra for the same element merges
they create a continuous spectrum
Process of absorption/emission spectrum creation
-electromagnetic radiation passes through a collection of atoms at ground state
-some radiation is absorbed and excites electrons to higher energy levels
-when this happens in the visible (n=2) region, a spectrometer analyzes the transmitted radiation to produce an absorption spectrum
-when electrons in the atom eventually return to their ground state, an emission spectrum is produced, as the energy absorbed is released
-these “jump and falls” are referred to as electron transitions
-the energy absorbed/released during these transitions are called “photons”
-each electron transition releases one photon of energy
Why are absorption/emission spectrum important
-shows that e- occupy distinct energy levels, allowing to quantitize electron transition
-each element has a unique spectra, and can be used like a barcode
-many transitions happen at the same time, but only those in the visible region produce a spectrum
(n=1=UV region, n=2=visible, n=3- IR region)
Limitations to the Bohr diagram model
electron transitions vary
spectral lines converge at higher energy levels
at n=infinity, ionization occurs where they overcome the attractive force of nucleus
electron repulsion must be overcome by opposite spin
electrons are both discrete particles and waves
Ionization energy calculation formula
IE = energy/mole of atom = Jmol-1 -> kJmol -1
(take energy value + multiply by avogadros #)
Limit of convergence =
ionization energy
define first IE
the minimum energy required to remove one mole of electrons from one mole of gaseous atoms in the ground state
X(g) -> 1e- + X 1+ (g)
what happens to ionization energy across a period, and down a group
IE increases as the nuclear charge increases. this is bc a higher nuclear charge = a stronger attraction of valence electrons, so it takes more energy to remove 1e-
IE decreases down a group, as the effect of adding energy levels (shielding effect) increases the distance b/w the nucleus and valence e-, lessening the nuclear pull/attraction
Which element has the lowest ionization energy
Fr (Francium)
What are the discontinuities in the 1st IE trend
Group 13 elements (s2 p1) vs Group 2 elements (s2) : group 13 elements have a lower 1st IE than group 2 elements due to the p orbital. This is bc a p orbital is further from the nucleus than an s orbital
Group 15 elements (s2p3) vs Group 16 elements (s2p4): group 16 elements have a lower 1st IE than group 15 elements bc of the repulsive effect of adding a paired electron to the p orbital
what is the increase in nuclear charge
the addition of p+ to the nucleus, which increases the attraction of valence e- to the nucleus
what is the shielding effect
refers to the effects of adding energy levels to atoms, which increases the distance between the nucleus and the valence electrons, which shields v e- from the nuclear pull
