f321 Flashcards

1
Q

State and explain the trend in first ionisation energies across a period.

A
  • First ionisation energies increase.
  • Outermost electrons in same shell.
  • Nuclear charge increases (more protons).
  • Atomic radius decreases (outermost electrons pulled closer to nucleus).
  • Nuclear attraction on outermost electron increases.
  • More energy needed to remove outermost electron.
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2
Q

State and explain the reactivity of elements down Groups 1&2.

A
  • First ionisation energy decreases down the group.
  • Reactivity increases.
  • Outermost electron is in a new shell further from the nucleus with more shielding effect from inner electrons.
  • Atomic radius increases (more shells).
  • Nuclear attraction on outermost electron decreases.
  • Although nuclear charge increases this is ‘far outweighed’ by more distance and shielding.
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3
Q

Explain why water is polar.

A
  • O is more electronegative than H.
  • Creates a permanent dipole across O-H bonds.
  • Water not symmetrical so dipoles do not cancel out.
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4
Q

State and explain the trend in atomic radii (size of atom) across a period.

A
  • Atomic radius (size of atom) decreases.
  • Outermost electron in the same shell.
  • Nuclear charge increases (more protons).
  • Nuclear attraction on electrons increases (electrons are pulled closer to the nucleus).
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5
Q

State and explain the trend in atomic radii (size of atom) down a group.

A
  • Atomic radii (size of atom) increases (more shells).
  • Greater distance of outermost electron to nucleus with more shielding from inner electrons.
  • Nuclear attraction on electrons decreases (electrons are not pulled as close to nucleus).
  • Although nuclear charge increases this is ‘far outweighed’ by more distance and shielding.
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6
Q

State and explain the reactivity of elements down Group 7.

A
  • Reactivity decreases.
  • More shells.
  • Greater distance to nucleus with more shielding from inner electrons.
  • Nuclear attraction (on electrons to be gained) decreases.
  • Atom is less likely to gain
    an electron to form the 1- ion.
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7
Q

Explain why there is an increase in melting/ boiling points from Na to Al.

A
  • Ionic charge increase: Na+, Mg2+ and Al3+ across period.
  • More delocalised electrons across period.
  • Attraction between positive metal ions and delocalised electrons (metallic bonding) increases.
  • More energy is needed to break the stronger metallic bonds.
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8
Q

State and explain the trend in melting/ boiling points down Group 7.

A
  • Melting/ boiling points increase.
  • More electrons down Group 7.
  • Strength of van der Waals’ forces between molecules increase.
  • More energy is needed to break the forces.
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9
Q

Describe the hydrogen bonding between two water molecules.

A
  • Well labelled diagram*
  • Water is polar
  • Permanent dipole across O-H bonds.
  • Hδ+ on one water molecule attracts lone pair of electrons on highly electronegative O atom of another water molecule.
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10
Q

Describe van der Waals’ (London) forces.

A
  • Occur in non-polar molecules (overall).
  • Temporary dipole forms in one molecule due to uneven distribution of electrons.
  • Induces a dipole in a neighbouring molecule.
  • Creates van der Waals’ forces between molecules.
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11
Q

In terms of structure and bonding, describe the melting/ boiling point, electrical conductivity and solubility of Diamond

A
  • Giant covalent.
  • High melting/ boiling point as the strong covalent bonds need a lot of energy to break.
  • Non-conductor of electricity as there are no charge carriers.
  • Insoluble in water as strong covalent bonds cannot be broken.
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12
Q

In terms of structure and bonding, describe the melting/ boiling point, electrical conductivity and solubility of Graphite

A
  • Giant covalent.
  • Weak van der Waals’ intermolecular forces between layers.
  • High melting/ boiling point as the strong covalent bonds need a lot of energy to break.
  • Good conductor of electricity as there are delocalised electrons free to move across the layers and carry charge.
  • Insoluble in water as strong covalent bonds cannot be broken.
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13
Q

In terms of structure and bonding, describe the melting/ boiling point, electrical conductivity and solubility of Sodium chloride

A
  • Giant ionic
  • High melting/ boiling point as the strong ionic bonds need a lot of energy to break.
  • Good conductor when aqueous/ molten as IONS are free to move and carry charge.
  • Non-conductor when solid as IONS are fixed in a lattice and cannot move or carry charge.
  • Soluble in water as polar H2O is attracted to Na+ and Cl- and ionic lattice breaks up.
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14
Q

In terms of structure and bonding, describe the melting/ boiling point, electrical conductivity and solubility of Iodine

A
  • Simple covalent
  • Weak van der Waals’ intermolecular forces between molecules.
  • Low melting/ boiling point as the weak intermolecular forces between molecules need less energy to break.
  • Non-conductor of electricity as there are no charge carriers.
  • Partially soluble in polar water.
  • Soluble in non-polar hexane (organic solvent) as hexane forms van der Waals’ with I2 molecules.
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15
Q

Outline how the ‘Electron Pair Repulsion Theory’ can be used to predict the shape of a molecule.

A
  • Electron pairs repel other electron pairs as far apart as possible.
  • Lone pairs of electrons repel more than bond pairs of electrons.
  • The number and type of electron pairs around the central atom determines the shape of a molecule.
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16
Q

State and explain two anomalous properties of H2O resulting from hydrogen bonding.

A
  • Ice is less dense than water: hydrogen bonds between the molecules are longer in ice than water holding H2O molecules further apart in an open lattice.
  • ‘Higher than expected’ melting/ boiling points: Hydrogen bonds between molecules are the strongest intermolecular force and therefore more energy is needed to break them.