ELECTRONIC STRUCTURE OF ATOMS Flashcards
Limitations of Bohr’s Theory (4)
- Bohr’s theory worked perfectly explaining the emmision spectrum of hydrogen, but failed to account for many lines in atoms with more than one e-
- Bohr’s theory does not account the wave like motion of e- so precise path of e- picture cannot be true
- Heisenberg’s Uncertainty Principle was in conflict with Borh’s theory
- Bohr’s theory did not explain splitted lines on emission spectrum so did not account for presence of sub-levels
ATOMIC ABSORPTION SPECTROMETRY = AAS
used to detect & calculate the concentration of an element present in a water sample.
A SPECTROSCOPE
used to analyse light emitted by elements.
PAULI EXCLUSION PRINCIPLE
no more than 2 electrons can occupy an orbital & they must have opposite spin.
FIRST IONISATION ENERGY of an atom
the minimum energy needed to remove the most loosely bound electron from a neutral gaseous atom in its ground state.
Na(g) → Na(g) + + e-
PRINCIPLE of AAS
- Atoms of an element in its ground state absorb light of a particular wavelength that is characteristic to that element.
- Amount of light absorbed is directly proportional to concentration of element present in sample
AUFBAU PRINCIPLE
e- will always occupy the lowesr available sub-level and energy level first
HUNDS RULE OF MAXIMUM MULTIPLICITY
when 2 or more orbitals of equal energy are available, electrons will occupy them singly before occupying them in pairs.
Erwin Schrodinger
Schrodinger’s equation
- Used to calculate the probability of finding an electron in a particular position in an atom.
- Led to defining the ATOMIC ORBITAL
ATOMIC ABSORPTION SPECTRUM
NOT AAS
series of dark lines on a coloured background.
Spectrometer
used to take measurements of the SPECTRA of colours produced by light
E2 - E1= hf
E2 - E1 = the difference in energy values between n=2 and the n=1 energy levels
h = Planck’s constant
f = frequency of light emitted
= = same/equal to
Nuclear charge
no. of protons (+) in nucleus attracting the electrons (-) towards it.
Screening effect of inner electrons
no. of full energy levels between the nucleus & outer electrons reduces the effect of nuclear charge.
ELECTRONEGATIVITY
relative attraction that an atom in a molecule has for a shared pair of electrons in a covalent bond.
Atomic Orbitals
region in space within an atom where there is a high probability of finding an electron
Sub-Level
subdivision of an energy level consisting of a group of atomic orbitals, all with the same energy
Werner Heisenberg…
Heisenberg’s Uncertainty Principle = it is impossible to measure at the same time the POSITION & the VELOCITY of an electron.
BALMER Series
lines in VISIBLE region of spectrum.
e- falling from higher energy levels back to n=2
Louis de Broglie …
Suggested all moving particles (electrons) have a wave like motion.
QUANTUM NUMBER
electrons in atoms can only have FIXED energy values
=QUANTISATION
Energy Level
a region of definate energy within an atom that electrons can occupy
The ATOMIC RADIUS of an element
half the distance between the nuclei of two atoms of the same element joined by a single covalent bond.
Measured using X-rays defraction.
Measured in nanometers (nm)
LINE SPECTRA of H as EVIDENCE of ENERGY LEVELS
- Electrons are restricted to a regions (orbits) with a FIXED energy value called ENERGY LEVEL .
- Each ENERGY LEVEL is identified by a PRINCIPAL QUANTUM NUMBER called “n” ….n=1 is lowest energy level, n=2 is next etc
- The energy of the electron in each energy level is QUANTISED (fixed @ a definite value….cannot have an energy value that lies between energy levels)
- An electron will occupy the LOWEST AVAILABLE energy level = it is in its GROUND STATE.
- When an electron absorbs energy (heat or electricity) it jumps from a lower energy level (E1) to a higher energy level (E2)= it is in a TEMPORARY (unstable) EXCITED STATE .
- When the electron falls from the higher energy level (E2) to a lower energy level (E1), a definite amount of energy is emitted equal to the difference in energy levels E2 - E1
–>This is seen as a PHOTON of LIGHT
-Amount of energy emitted is represented by
E2 - E1= hf
