ELECTRONIC STRUCTURE OF ATOMS

Question 1

Limitations of Bohr’s Theory (4)

Answer 1

• 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

Question 2

ATOMIC ABSORPTION SPECTROMETRY = AAS

Answer 2

used to detect & calculate the concentration of an element present in a water sample.

Question 3

A SPECTROSCOPE

Answer 3

used to analyse light emitted by elements.

Question 4

PAULI EXCLUSION PRINCIPLE

Answer 4

no more than 2 electrons can occupy an orbital & they must have opposite spin.

Question 5

FIRST IONISATION ENERGY of an atom

Answer 5

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-

Question 6

PRINCIPLE of AAS

Answer 6

• 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

Question 7

AUFBAU PRINCIPLE

Answer 7

e- will always occupy the lowesr available sub-level and energy level first

Question 8

HUNDS RULE OF MAXIMUM MULTIPLICITY

Answer 8

when 2 or more orbitals of equal energy are available, electrons will occupy them singly before occupying them in pairs.

Question 9

Erwin Schrodinger

Answer 9

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

Question 10

ATOMIC ABSORPTION SPECTRUM

NOT AAS

Answer 10

series of dark lines on a coloured background.

Question 11

Spectrometer

Answer 11

used to take measurements of the SPECTRA of colours produced by light

Question 12

E2 - E1= hf

Answer 12

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

Question 13

Nuclear charge

Answer 13

no. of protons (+) in nucleus attracting the electrons (-) towards it.

Question 14

Screening effect of inner electrons

Answer 14

no. of full energy levels between the nucleus & outer electrons reduces the effect of nuclear charge.

Question 15

ELECTRONEGATIVITY

Answer 15

relative attraction that an atom in a molecule has for a shared pair of electrons in a covalent bond.

Question 16

Atomic Orbitals

Answer 16

region in space within an atom where there is a high probability of finding an electron

Question 17

Sub-Level

Answer 17

subdivision of an energy level consisting of a group of atomic orbitals, all with the same energy

Question 18

Werner Heisenberg…

Answer 18

Heisenberg’s Uncertainty Principle = it is impossible to measure at the same time the POSITION & the VELOCITY of an electron.

Question 19

BALMER Series

Answer 19

lines in VISIBLE region of spectrum.

e- falling from higher energy levels back to n=2

Question 20

Louis de Broglie …

Answer 20

Suggested all moving particles (electrons) have a wave like motion.

Question 21

QUANTUM NUMBER

Answer 21

electrons in atoms can only have FIXED energy values

=QUANTISATION

Question 22

Energy Level

Answer 22

a region of definate energy within an atom that electrons can occupy

Question 23

The ATOMIC RADIUS of an element

Answer 23

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)

Question 24

LINE SPECTRA of H as EVIDENCE of ENERGY LEVELS

Answer 24

• 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