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Physical Chem AS Mock > Periodicity > Flashcards

Periodicity Flashcards

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1
Q

Explain why xenon has a lower first ionisation energy than neon. (3)

A
  • more shielding
  • bigger atomic radius
  • less nuclear attraction
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2
Q

Explain why the first ionisation energies show a general increase from Li to Ne. (3)

A
  • atomic radii increases
  • nuclear charge increases
  • shielding remains the same
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3
Q

Explain the difference between the first ionisation energies of Li and Na. (3)

A
  • atomic radii increases
  • more shielding
  • distance outweighs nuclear charge
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4
Q

Explain why first ionisation energies show a general increase across Period 3, Na - Ar. (3)

A
  • atomic radii decreases
  • number of protons in nucleus increases
  • nuclear attraction increases
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5
Q

What is the equation, including state symbols, for the third ionisation energy of Sodium? (1)

A

Na2+ (g) –> Na3+ (g) + e-

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6
Q

Explain why less energy is needed to ionise gaseous atoms of rubidium than gaseous atoms of sodium. (3)

A
  • Rb = more shells
    ^ more shielding
    ^ less attraction
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7
Q

Explain whether a Barium ion is larger, smaller or the same size as a Barium atom. (2)

A
  • smaller
  • less shielding
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8
Q

Define: first ionisation energy

A

the energy change when each atom in 1 mole of gaseous atoms loses an electron

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9
Q

Explain why the first ionisation energy of B is less than that of Be. (2)

A
  • In B, the electron is being removed at a higher energy
  • An s-orbital electron is lost in Be & a p-orbital electron is lost in B
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10
Q

Explain why less energy is needed to remove an electron from a radium atom than from a calcium atom. (3)

A
  • atomic radii of radium is bigger
  • radium has more shells
    ^ more shielding
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11
Q

Explain why a nitrogen atom is larger than an oxygen atom. (4)

A
  • Nitrogen has less protons
  • similar shielding
  • Nitrogen has weaker nuclear attraction
  • Shell drawn in less by nuclear charge in Nitrogen
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12
Q

What determines the order of elements in The Periodic Table?

A

the number of protons in the nucleus

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13
Q

Why does phosphorus have a larger first ionisation energy than sulfur? (1)

A
  • P atoms have less repulsion between p-orbital electrons than S atoms
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14
Q

Define: periodicity

A
  • trend in the properties of elements across a period repeated in the next row
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15
Q

Define: metallic bonding

A

electrostatic force of attraction between metal ions and delocalised electrons

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16
Q

State the element in Period 3 with the highest first ionisation energy

A

Silicon

17
Q

Describe the structure and bonding present in silicon. (2)

A
  • giant macromolecular structure
  • covalent bonds between silicon atoms
18
Q

Explain why aluminium has a higher melting point than sodium. (3)

A
  • Aluminium has a higher charge
    ^ more delocalised electrons
  • stronger metallic bonding
19
Q

Explain why aluminium does not follow the general trend in first ionisation energies across the period. (3)

A
  • lower than Mg
  • In Al, electron removed from p-subshell
  • electron is of higher energy
20
Q

Describe a similarity between copper and graphite. (1)

A
  • both have layers of atoms that can slide over one another
21
Q

Explain why copper has a high melting point. (2)

A
  • contains positive metal ions
    ^ strong attraction b/w them
22
Q

Identify the Period 3 element with the lowest melting point and explain why. (3)

A
  • Argon
  • simple molecular structure
  • molecules held together by weak intermolecular forces which require very little energy to overcome
23
Q

Explain why the second ionisation energy of Aluminium is larger than the first. (1)

A
  • electron is being removed from an s-orbital rather than a p-orbital
24
Q

Describe the structure & bonding of graphite + why it has such a high melting point. (5)

A
  • arranged in layers
  • Carbon atoms in each layer connected by 3 covalent bonds
  • weak intermolecular forces b/w layers

high m.p: - millions of strong covalent bonds
^ require a large amount of energy to overcome

25
Q

Explain why the third ionisation energy of Silicon is higher than its second. (2)

A
  • Si2+ ions are smaller than Si+ ions
  • more energy required
26
Q

Explain why an oxygen atom is smaller than a nitrogen atom. (4)

A
  • oxygen = more protons
  • similar shielding
  • stronger attraction between protons & electrons
  • electrons drawn in more by increased nuclear charge
27
Q

Explain why silicon has the highest melting point in Period 3. (3)

A
  • giant macromolecular structure
  • held together by strong covalent bonds
  • high amount of energy required to overcome them —> high m.p
28
Q

Explain why sulfur has a higher melting point than phosphorus. (3)

A
  • both have weak intermolecular forces
  • strength of induced dipole-dipole forces increases with the electron number
  • sulfur has more electrons
29
Q

Explain the trend in melting point across the elements of Period 2. (6)

A

Li to Be: metallic bonding
^ gets stronger as number of delocalised electrons increases

B & C: giant structure of atoms
^ high amount of energy needed to break strong covalent bonds

N to Ne: simple molecules
^ weak induced dipole-dipole forces between molecules

30
Q

Explain why strontium has a lower second ionisation energy than magnesium. (3)

A
  • Sr+ = larger ionic radius
    ^ experiences more shielding
  • less nuclear attraction to outermost electrons
31
Q

Describe the bonding and structure of Magnesium. (2)

A
  • positive ions
  • surrounded by delocalised electrons
32
Q

Explain why magnesium has a higher melting point than sodium. (2)

A
  • Mg+ has more protons
  • attracts delocalised electrons more strongly
33
Q

Does Mg2+ or Na+ have a smaller ionic radius? Explain why. (2)

A
  • Mg2+
  • bcs higher nuclear charge & similar shielding

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