Bonding (old) Flashcards

1
Q

Description of ionic bonds

A
  • metal and non metal
  • transfer electrons
  • opposite charged ions attracted by electrostatic forces
  • in a giant ionic lattice structure
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2
Q

Properties of ionic compounds

A
  • crystalline solid at room temp
  • high m.p
  • only conduct electricity when motlen or dissolved, ions now free to move to carry a current throughout the structure
  • brittle and shatter easily
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3
Q

Why are ionic compounds brittle?

A
  • because they form a lattice of alternating +/- ions, a small displacement move the ions and causes contact between ions with same charges
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4
Q

Why does ionic compounds have high melting point?

A
  • need a lot of energy to break the giant lattice of ions with strong electrostatic forces between + - charged ions
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5
Q

Description of molecular covalent bonding

A
  • in a simple molecular structure
  • non metal non metal
  • shared pair of electrons
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6
Q

Properties of molecular covalent bonding

A
  • gas and liquid in room temp.
  • low m.p.
  • poor solubility
  • don’t conduct electricity- not charged so no free ions, no delocalised electrons
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7
Q

Why is molecular Covalent bond has low m.p?

A
  • weak intermolecular forces between MOLECULES so easy to break, less energy needed to overcome
  • Types of forces : van der waals, permanent dipoles, hydrogen bonds
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8
Q

Why is molecular covalent has low solubility ?

A
  • no charged particles
  • don’t interact with polar water molecules
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9
Q

What affects the strength of covalent bond ?

A
  • atomic radius- the bigger the atom is, the greater distance so weaker attraction between nucleus and electrons
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10
Q

Factors affect strength of ionic bonds (2)

A
  • charge of ions- higher charge( more protons +) = stronger attraction of nucleus and e-
  • size of ions - smaller ion( smaller atomic radius ) = less distance = stronger attraction of nucleus and e-
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11
Q

Description of macromolecular covalent bonding

A
  • non metals x2
  • shared pair of electrons
  • macromolecular structure
    E.g. diamond, graphite, silicon dioxide, silicon
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12
Q

Properties of macromolecular covalent bond

A
  • solid at room temp.
  • high m.p.
  • insoluble
  • diamond and sand don’t conduct electricity, graphite does as free delocalised electrons between layers
  • don’t conduct when molten
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13
Q

Why is m.p high in macro.cov bond

A
  • many strong covalent bonds in macro structure = need a lot of energy to break
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14
Q

Bonding + structure of diamond

A
  • each C atom bond with 4 other tetrahedral arrangement
  • does not conduct electricity but good thermal conductor
  • very high m.p.
  • extremely hard
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15
Q

Bonding + structure of graphite

A
  • each C atom bond with 3 other C atoms , the 4th e- is delocalised
  • hexagonal arrangement
  • layered structure
  • weak intermolecular forces between layers
  • high m.p.
  • insoluble, covalent bonds too strong to break
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16
Q

Why graphite is soft ?

A
  • weak intermolecular force between layers = can slide over each other
17
Q

why is graphite light weight ?

A
  • layers are far apart compared length of covalent bonds= low density
18
Q

Bonding and structure of metallic bond

A
  • solid at room temp.
  • metal x2
  • giant metallic lattice structure
  • electrostatic force of attraction between metal + ions / delocalised e-
19
Q

Properties of metallic bond

A
  • high m.p
  • insoluble- as metallic bonds are non polar
  • good conductor of electricity- delocalised electrons give charge to the whole structure
  • malleable
20
Q

Why is metallic bond high m.p.

A
  • strong electrostatic forces between + ions and delocalised sea of e-
21
Q

Factors affect metallic bond

A
  • size of ions- larger ions = larger ionic radius = weaker attraction of positive ions to delocalised e- = decrease strength
  • charge of ions- higher charge = stronger attraction of positive ions to delocalised e- = greater strength
22
Q

Why metals are malleable ?

A
  • layers of ion can slide over each other without any disruption to metallic bond = + ions still attract to delocalised e-
23
Q

Why is alloy harder than pure metals ?

A
  • a mixture of metals= different sized atoms distort regular layered structure so layers cannot slide = harder