TOPIC 2: ATOMIC STRUCTURE Flashcards
Where’s the bulk of the mass of an atom?
Almost the whole mass of the atom is concentrated in the nucleus,
THOUGH the nucleus takes up an extremely small space relative to the size of the whole atom
Fun fact! Size of an atom is easily more than 10,000 times that of the nucleus
Representation of element
Nucleon number = Mass number
Proton number = Atomic number
2 Main Properties of Isotopes
Similar chemical properties: Same number and arrangement of electrons
Different physical properties: Different number of neutrons, different masses
Factors determining angle of deflection
angle of deflection ∝ |q/m| [for magnitude only]
If want find direction as well, remove modulus sign
Things to note when drawing electric field
- Before entering field, straight line
- In the field, start bending
- Once exit field, back to straight line
*Electron’s angle of deflection should be larger than proton’s
Overview of arrangement of electrons
Electrons –> Orbital (2 e-) –> Subshell –> Principal Quantum Shell
s subshell: 1 orbital
p subshell: 3 orbitals
d subshell: 5 orbitals
f subshell: 7 orbitals
Shape of s orbital
1 spherical orbital (NOT circular)
Shape of p orbital
3 dumbbell shaped orbitals (along x, y, z axis respectively)
- Putting all 3 together is a sphere, so each p orbital is 1/3 the volume of a sphere
- All 3 p orbitals are degenerate (aka same energy)
Shape of d orbital
5 dumbbell shaped orbitals
* between xy / xz / yz & along x²-y² / z²
- Putting all 5 together is a sphere, so each d orbital is 1/5 the volume of a sphere
- All 5 d orbitals are degenerate (aka same energy)
Energy of electrons in the shells
Further away from nucleus –> Less strongly attracted –> Higher energy level
Subshell: s < p < d < f
Principal Quantum shell: n increase, energy increase
** 4s < 3d
Rules for arranging electrons
Pauli’s Exclusion Principle: Same orbital must have opposite spins
Hund’s Rule: Orbitals of a subshell must be singly occupied first with parallel spins before pairing occurs
The Aufbau Principle: Place electrons into orbitals, starting with those of the lowest energy and then working upwards
Filling / Removing electrons from 3d and 4s subshells
Fill 4s then 3d (lower energy, then higher energy)
Remove 4s then 3d (higher energy, then lower energy)
* Becomes higher energy cos of shielding effect (repulsion with inner electrons)
Ways to present electronic configurations (5)
Draw e.c: (draw parallel lines)
Write/state full e.c: 1s²2s²……
Write/state short form notation of e.c: [Ar]4s¹
Draw energy level diagram: (vertical ‘Energy’ axis)
Describe subshell using clearly labelled diagrams (draw out)
*ec : electronic configuration
Period / Group number
Period number indicates the principal quantum number of the valence shell
Group number indicates the number of electrons in the valence shell
ONLY Anomalous Electronic Configurations
Cr: [Ar] 3d⁵4s¹ (NOT 3d⁴4s²)
Cu: [Ar] 3d¹⁰4s¹ (NOT 3d⁹4s²)
*Energetically preferred configuartions due to the symmetrical 3d electron cloud around the nucleus, so greater stability
Factors affecting electrostatic forces of attraction (EFA)
- Effective Nuclear Charge (Nuclear Charge - Shielding Effect)
Nuclear Charge
* No. of protons in the nucleus of atom
Shielding Effect
* No. of inner shell electrons
- Distance of valence electron from nucleus
* Valence electrons located in a shell with a larger principal quantum number
* Valence electrons further away from nucleus
Shielding (or screening) effect
Shielding (or screening) effect refers to the inter-electronic repulsion between the valence electrons and the inner shell electrons
[Definition] Ionsiation Energy
The first ionisation energy is the energy required to remove 1 mole of electrons from 1 mole of gaseous atoms to form 1 mole of singly charged gaseous cations.
The second ionisation energy is the energy required to remove 1 mole of electrons from 1 mole of singly charged gaseous cations to form 1 mole of doubly charged gaseous cations.
Trend of First Ionisation Energy (across period)
- Protons increase (nuclear charge increase)
- Inner shell electrons remain the same (approximately constant shielding effect)
- Increase in effective nuclear charge
- Increase in electrostatic forces of attraction
- More energy needed to remove valence electron
- Increase in Ionisation Energy
EXCEPTION:
1. Decrease from grp 2 to 13
* electron in p subshell to be removed is of a higher energy level than the electron in s subshell to be removed
* Less energy required to remove valence electron
2. Decrease from group 15 to 16
* Inter-electronic repulsion between the paired electrons in the same orbital
* Less energy required to remove valence electron
Trend of First Ionisation Energy (down group)
- Both protons and inner shell electrons increase
- Both nuclear charge and shielding effect increase
- Valence electrons in a shell with a larger principal quantum number, so further away from nucleus
- Weaker electrostatic forces of attraction
- Less energy needed to remove valence electron
- Decrease in Ionisation Energy
Trend of successive Ionisation Energy
- Proton remain the same, so nuclear charge constant
- Shielding effect decrease as electrons are removed
- Effective nuclear charge increases
- Stronger electrostatic forces of attraction
- More energy to remove valence electrons
- Increase in Ionisation Energy
Deducing electronic configuration using successive Ionisation Energy graph
Sudden increase in ionisation energy –> Change in principal quantum shell –> Indicates number of valence electrons
[Definition] Atomic Radius
Atomic radius is defined as half the distance between the centres of two adjacent atoms found in the structure of the element.
[Definition] Ionic Radius
Ionic radius is defined as the average distance between the nuclei of two adjacent ions of the same element.
[Definition] Electronegativity
Electronegativity refers to the ability of an atom to attract the shared pair(s) of electrons in a covalent bond.
Trend of atomic / ionic radii across period
- Protons increase (nuclear charge increase)
- Inner shell electrons same (shielding effect approximately constant)
- Effective nuclear charge increase
- Increase in electrostatic forces of attraction
- Atomic / cationic / anionic radii decreases
Trend of atomic / ionic radii down group
- Both protons and inner shell electrons increase
- Both nuclear charge and shielding effect increase
- Valence electrons in a shell with a larger principal quantum number, so further away from nucleus
- Decrease in electrostatic forces of attraction
- Atomic / cationic / anionic radii increases
Trend of electronegativity across period
- Proton increase (nuclear charge increase)
- Inner shell electrons same (shielding effect approximately constant)
- Effective nuclear charge increase
- Electrostatic forces of attraction increases
- Electronegativity increases
Trend of electronegativity down group
- Both protons and inner shell electrons increase
- Both nuclear charge and shielding effect increase
- Valence electrons in a shell with a larger principal quantum number, so further away from nucleus
- Decrease in electrostatic forces of attraction
- Decrease in electronegativity
Cationic / Anionic Radii
Cations < corresponding atom
Anions > corresponding atom
Anions > Cations
