Atomic Structure Flashcards

1
Q

Protons

Relative mass: ?

Relative charge: ?

Behaviour in electric field: ?

A
  • Relative mass: 1
  • Relative charge: +1
  • Behaviour in electric field: deflected to negative plate (smaller deflection than electron due to greater mass)
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2
Q

Neutrons

Relative mass: ?

Relative charge: ?

Behaviour in electric field: ?

A
  • Relative mass: 1
  • Relative charge: 0
  • Behaviour in electric field: no deflection as no charge
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3
Q

Electrons

Relative mass: ?

Relative charge: ?

Behaviour in electric field: ?

A
  • Relative mass: 1/1836
  • Relative charge: -1
  • Behaviour in electric field: deflection to positive plate (large deflection due to small mass)
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4
Q

Define

isotopes

A

Atoms of the same element/with the same atomic number/with the same number of protons but different masses/mass numbers/number of neutrons.

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5
Q

What are principal energy levels/principal quantum shells?

A

These are regions at various distances from the nucleus that may contain up to a certain number of electrons (with similar energies and similar distances from the nucleus).

As the shell number increases, the energy of the electrons increases.

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6
Q

What is a subshell?

A

A group of orbitals with the same energy level that differ in their orientation in space.

Each principal quantum shell contains a different number of subshells.

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7
Q

What is an orbital?

A

An atomic orbital is a region of space around the nucleus where there is a maximum probabilty of electrons to be found.

Each orbital can contain up to two electrons. For orbitals with two electrons, the electrons have opposite spins two minimise repulsion.

When e-s are placed in a subshell of orbitals of equal energy, they occupy them singly before pairing takes place.

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8
Q

State the number of orbitals and the maximum number of electrons in these subshells:

s: ?
p: ?
d: ?

A
  • s: 1 orbital, 2 electrons
  • p: 3 orbitals, 6 electrons
  • d: 5 orbitals, 10 electrons
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9
Q

What are the shapes of the s and p orbitals?

A
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10
Q

In what order are atomic orbitals filled?

A

Electrons will always occupy the lowest available energy level orbital first (Aufbau’s principle)

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11
Q

State the exceptions to the expected pattern of electronic configuration

A

Chromium:

1s2 2s2 2p6 3s2 3p6 4s1 3d5

Copper:

1s2 2s2 2p6 3s2 3p6 4s1 3d10

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12
Q

Define

ionisation energy

A

The minimum energy needed to remove 1 mole of e-s from 1 mole of gaseous atoms to form 1 mole of unipositive ions.

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13
Q

What 4 factors influence ionisation energy?

A
  1. Distance from nucleus (atomic radius)
    • farther away means lower attraction
  2. Nuclear charge
    • as nuclear charge increases, attraction increases (usually leading to a smaller radius, e.g. between isoelectronic ions of different elements)
  3. Shielding
    • outer electrons are pushed away by full subshells of inner electrons
  4. Spin-pair / Inter-electron repulsion
    • an electronic configuration having two electrons in an orbital is less stable than having one
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14
Q

What trend in ionisation energy can be seen down a group?

A

Ionisation energy decreases down a group

  • No. of shells increases, so atomic radius increases
  • Larger number of shells means more inner electrons, so shielding increases
  • Higher number of protons means higher attraction, but this factor is outweighed by shielding and distance from nucleus
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15
Q

What trend in ionisation energy can be seen across a period?

A

Ionisation energy generally increases across a group

  • Number of protons increases so positive nuclear charge and attraction increase
  • No. of shells and shielding remain reasonably constant
  • Higher attraction leads to smaller atomic radius
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16
Q

Explain the first drop in ionisation energy when moving across periods 2 or 3.

(i.e. between Be and B or Mg and Al)

A

The electron to be removed is in a p subshell rather than an s subshell, and so is slightly further from the nucleus and has an additional subshell of inner electrons shielding it.

17
Q

Explain the second drop in ionisation energy when moving across periods 2 or 3.

(i.e. between N and O or P and S)

A

The previous element has half-subshell stability. The spin-pair repulsion in O and S leads to a less stable configuration, so the electron is more eaily removed.

18
Q

How can you interpret this information from a successive ionisation energy graph?

  • Total no. of electrons
  • No. of shells occupied
  • No. of subshells occupied
  • No. of electrons in each shell
  • No. of electrons in each subshell
A
  • Total no. of electrons: no. of plot points on graph
  • No. of shells occupied: no. of sections between large jumps in energy
  • No. of subshells occupied: no. of section between small jumps in energy
  • No. of electrons in each shell: no. of electrons between two large jumps in energy
  • No. of electrons in each subshell: no. of electrons between small jumps in energy
19
Q

Describe how ionic radius differs between positive and negative ions.

A

Ionic radius: describes the size of an ion

  • Positive ion: smaller radius than original neutral atom because shell no. decreases, screening effect decreases but the attraction of nucleus increases.
  • Negative ion: larger ionic radius than neutral atom because e-s added while nuclear charge remains same
  • Negative ions always larger than positive ions in the same period as they have one more shell