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EJC CHEMISTRY > TOPIC 2: ATOMIC STRUCTURE > Flashcards

TOPIC 2: ATOMIC STRUCTURE Flashcards

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1
Q

Where’s the bulk of the mass of an atom?

A

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

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2
Q

Representation of element

A

Nucleon number = Mass number

Proton number = Atomic number

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3
Q

2 Main Properties of Isotopes

A

Similar chemical properties: Same number and arrangement of electrons

Different physical properties: Different number of neutrons, different masses

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4
Q

Factors determining angle of deflection

A

angle of deflection ∝ |q/m| [for magnitude only]

If want find direction as well, remove modulus sign

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5
Q

Things to note when drawing electric field

A
  1. Before entering field, straight line
  2. In the field, start bending
  3. Once exit field, back to straight line

*Electron’s angle of deflection should be larger than proton’s

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6
Q

Overview of arrangement of electrons

A

Electrons –> Orbital (2 e-) –> Subshell –> Principal Quantum Shell

s subshell: 1 orbital
p subshell: 3 orbitals
d subshell: 5 orbitals
f subshell: 7 orbitals

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7
Q

Shape of s orbital

A

1 spherical orbital (NOT circular)

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8
Q

Shape of p orbital

A

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)
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9
Q

Shape of d orbital

A

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)
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10
Q

Energy of electrons in the shells

A

Further away from nucleus –> Less strongly attracted –> Higher energy level

Subshell: s < p < d < f
Principal Quantum shell: n increase, energy increase

** 4s < 3d

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11
Q

Rules for arranging electrons

A

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

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12
Q

Filling / Removing electrons from 3d and 4s subshells

A

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)

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13
Q

Ways to present electronic configurations (5)

A

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

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14
Q

Period / Group number

A

Period number indicates the principal quantum number of the valence shell

Group number indicates the number of electrons in the valence shell

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15
Q

ONLY Anomalous Electronic Configurations

A

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

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16
Q

Factors affecting electrostatic forces of attraction (EFA)

A
  1. Effective Nuclear Charge (Nuclear Charge - Shielding Effect)

Nuclear Charge
* No. of protons in the nucleus of atom

Shielding Effect
* No. of inner shell electrons

  1. Distance of valence electron from nucleus
    * Valence electrons located in a shell with a larger principal quantum number
    * Valence electrons further away from nucleus
17
Q

Shielding (or screening) effect

A

Shielding (or screening) effect refers to the inter-electronic repulsion between the valence electrons and the inner shell electrons

18
Q

[Definition] Ionsiation Energy

A

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.

19
Q

Trend of First Ionisation Energy (across period)

A
  • 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

20
Q

Trend of First Ionisation Energy (down group)

A
  • 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
21
Q

Trend of successive Ionisation Energy

A
  • 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
22
Q

Deducing electronic configuration using successive Ionisation Energy graph

A

Sudden increase in ionisation energy –> Change in principal quantum shell –> Indicates number of valence electrons

23
Q

[Definition] Atomic Radius

A

Atomic radius is defined as half the distance between the centres of two adjacent atoms found in the structure of the element.

24
Q

[Definition] Ionic Radius

A

Ionic radius is defined as the average distance between the nuclei of two adjacent ions of the same element.

25
Q

[Definition] Electronegativity

A

Electronegativity refers to the ability of an atom to attract the shared pair(s) of electrons in a covalent bond.

26
Q

Trend of atomic / ionic radii across period

A
  • 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
27
Q

Trend of atomic / ionic radii down group

A
  • 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
28
Q

Trend of electronegativity across period

A
  • Proton increase (nuclear charge increase)
  • Inner shell electrons same (shielding effect approximately constant)
  • Effective nuclear charge increase
  • Electrostatic forces of attraction increases
  • Electronegativity increases
29
Q

Trend of electronegativity down group

A
  • 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
30
Q

Cationic / Anionic Radii

A

Cations < corresponding atom
Anions > corresponding atom
Anions > Cations

EJC CHEMISTRY (9 decks)

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