Atomic Structure Flashcards
Orbital
a region of space in which up to two electrons are likely to be located
Subshell
A subdivision of an electron shell, containing a fixed number of orbitals at
the same energy level
Electron shell
An energy level within an atom that may be occupied by a fixed number of electrons
how many electrons in each subshell
S, P, D, F
S- 2
P- 6
D- 10
F- 14
The Pauli exclusion principle
Pauli’s Exclusion Principle states that only a maximum of two electrons may be found in a given atomic orbital, and a filled orbital’s electrons will have opposite spin.
The Aufbau principle
The Aufbau principle states that subshells are filled in order from lowest energy to highest energy, and that a lower-energy subshell will be completely filled before electrons move into a higher-energy subshell.
Hund’s rule of maximum multiplicity
‘electrons in a partially filled subshell will arrange themselves so as to form the maximum number of half-filled orbitals’.
meaning every orbital in a subshell being half filled before any orbital in that subshell is doubly occupied, minimising repulsion
Exceptions to the Aufbau principle
Copper (Cu) 4s^1 3d^10
Chromium (Cr) 4s^1 3d^5
Atomic radius
half the distance between two nuclei of a diatomic molecule, assuming a single covalent bond between two identical atoms.
Core charge
a measure of the net attractive force felt by the valence shell electrons towards the nucleus.
How to calculate core charge?
number of protons in nucleus
-
number of electrons in inner shells
Electronegativity
the strength with which atoms of an element attract electrons when they are chemically combined with another element
trend in Atonic Radii when moving Down a Group
As you move down a group the atomic radius
increases
EXPLANATION:
• the number of occupied energy levels increases
• core charge remains constant
trend in Atomic Radii across a Period
As you move across a period the atomic radius decreases.
EXPLANATION:
• The number of occupied energy levels remains constant
• Core charge increases resulting in the valence electrons being more strongly attracted to the nucleus.
Trend in electro negativity down a group
As you move down a group the electronegativity decreases
EXPLANATION:
• Number of occupied energy levels increases, therefore the atomic radius increases
• Nuclear charge increases, however electron shielding also increases so the core charge remains the same overall
trend in electronegativity across a period
As you move across a period the electronegativity increases
EXPLANATION
• Number of occupied energy levels stays constant, and the atomic radius decreases (due to increasing core charge)
• Nuclear charge increases but number of inner shell electrons stays the same (shielding remains constant) so the core charge increases
What causes a spectral line on an emission spectrum?
The relapse of an excited electron from a higher to a lower shell causes the release of a photon of light energy with a specific wavelength.
What causes an electron to be promoted to a higher energy level?
The absorption of energy
First ionisation energy
• The first ionisation energy is defined as the amount of energy required to remove an electron from each of a mole of gaseous atoms.
(Produces gaseous ions with +1 charge)
Trend in ionisation energy down a group of the Periodic Table
As you move down a group the first ionisation energy decreases
EXPLANATION:
• the number of occupied energy levels increases so the atomic radius increases (and therefore nuclear attraction decreases)
• Despite the nuclear charge increasing, electron shielding also increases so core charge remains constant
• Therefore less energy is required to remove an electron.
Trend in ionisation energy across a period of the Periodic Table
As you move across a period the first ionisation energy increases.
EXPLANATION:
• The number of occupied energy levels remains constant, and the atomic radius decreases so the outermost electrons are closer to the nucleus.
• Core charge increases resulting in the outermost electrons being more strongly attracted to the nucleus.
• Therefore more energy is required to remove an electron.
Trend of successive ionisation energies
as the cation gets a greater positive charge, it becomes progressively more difficult to remove a negatively charged electron.
Difference between the reactions of metals and non-metals
When metals react they lose electrons from their valence shell. More reactive metals lose electrons more easily.
When non-metals react they accept electrons into their valence shell. More reactive non-metals accept electrons more easily
Reactivity of metals across period
reactivity of the metals decreases across a period
EXPLANATION:
When metals react they lose electrons (form positive ions). As we move across a period
- Core charge increases and electron shielding remains constant
This results in a general increase in the energy required to remove an electron
Reactivity of metals down a group
Reactivity of metals increases down a group
EXPLANATION:
-Core charge remains constant
-number of energy levels and electron shielding increases
-general decrease in energy required to remove an electron
Reactivity of non metals across period
Reactivity increases across period
EXPLANATION:
When non metals react they gain electrons (form anions). As we move across a period
- Core charge increases and electron shielding remains constant
- non metals with greatest attractive force will be the most reactive.
Reactivity of non metals down a group
Reactivity decreases down a group
EXPLANATION:
- Core charge remains constant
- number of energy levels and electron shielding increases
- general decrease in attractive force
Isotope
Isotopes of an element have the same atomic number (number of protons) but a different mass number (number of protons and neutrons).
Relative atomic mass (RAM) (Ar)
The relative atomic mass is the average mass of all of the isotopes of an element weighted for their relative abundance.
(relative abundance * isotopic mass) +
(relative abundance * isotopic mass)/100
What does a mass spectrometer measure?
an accurate mass for the isotopes of an element, and the relative abundance of those isotopes
