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Chemistry > Chemical Bonding > Flashcards

Chemical Bonding Flashcards

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1
Q

Ionic bonding

A

The electrostatic attraction between oppositely charged ions

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2
Q

Giant ionic lattice

A

A three dimensional structure of oppositely charged ions held together by ionic bonds

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3
Q

Properties of ionic compounds

A
  • High melting point and boiling point as ionic bonds are strong and have to all be broken to change states
  • Conducts electricity when molten or aqueous as ions are free to move to carry charge
  • Dissolves in polar substances. Six O- rip off a cation while 4 H+ rip off the anion
  • Don’t dissolve in non-polar solvents
  • Don’t conduct electricity as a solid
  • Tend to be brittle as when the layers shift, the opposite ions will repel
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4
Q

Isoelectronic

A

To have the same number of electrons e.g. Si4+ is isoelectronic with Ne

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5
Q

Covalent bond

A

The electrostatic attraction between two nuclei and the same shared pair of electrons

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6
Q

Lone pair

A

An outer shell pair of electrons that is not involved in chemical bonding

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7
Q

The octet rule

A

The rule that only 8 electrons occupy one shell when bonding with other atoms

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8
Q

Dative covalent bond

A

A covalent bond formed when a lone pair has been provided by one of the bonding atoms

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9
Q

Properties of diamond

A
  • Doesn’t conduct electricity as there are no free charge carriers available
  • Giant covalent lattice
  • Tetrahedral C arrangement
  • Bare hard as covalent bonds are strong and a lot of energy is needed to break them all
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10
Q

Properties of graphite

A
  • Giant covalent structure
  • Conducts electricity as there is a layer of delocalised electrons between the C layers that can carry charge
  • C arranged in layers
  • Layers can slide over each other so graphite can acts as a lubricant
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11
Q

Allotrope

A

Refers to one or more forms of an elementary substance

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12
Q

Metallic bonding

A

The electrostatic attraction between cations and a sea of delocalised electrons

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13
Q

Properties of metals

A
  • High melting and boiling points as there is a large electrostatic attraction between the cations and the delocalised electrons
  • Good conductors as delocalised electrons can carry charge
  • Even better conductors when molten as the cations can carry charge too
  • Insoluble in water or non-polar substances as the metallic bonds are too strong relative to intermolecular forces
  • Malleable
  • Ductile
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14
Q

Trend in mp/bp of metals (I,II,III)

A
  • Increases along the group as the number of electrons are increasing and the charge on the cations are too
  • Decreases down a group as the outer electron is far from the nucleus and can be removed easily
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15
Q

Malleability

A

The ability to change shape without breaking

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16
Q

Ductility

A

The ability to be pulled into a thin wire

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17
Q

Alloys

A

A mixture of metals; not a compound
E.g. stainless steel = steel + chromium
Brass = copper + zinc

18
Q

Use of alloys

A

Alloys make metals harder as the ions sizes of the metals are different, so the layers cannot slide over each other

19
Q

Valence Shell Electron Pair Repulsion theory

A

Each bonded pair repel themselves so they are as far away from eachother
Lone pairs repel more than bonding pairs

20
Q

Linear

A

e.g. BeCl2
2 bonding pairs 0 lone pairs
bond angle 180 degrees

21
Q

Trigonal planar

A

e.g. BF3
3 bonding pairs 0 lone pairs
bond angle 120 degrees

22
Q

Tetrahedral

A

e.g. CH4
4 bonding pairs 0 lone pairs
bond angle 109.5 degrees

23
Q

Trigonal pyramid

A

e.g. NH3
3 bonding pairs 1 lone pair
bond angle 107 degrees

24
Q

Bent

A

e.g. H2O
2 bonding pairs and 2 lone pairs
bond angle 104.5 degrees

25
Q

Octahedral

A

e.g. SF6
6 bonding pairs 0 lone pairs
bond angle 90 degrees

26
Q

The three types of intermolecular force

A

van der Waals’
Permanent dipole-dipole
Hydrogen bonding

27
Q

How do van der Waals’ forces arise?

A

Electrons are constantly moving. At any instant, the distribution may not be symmetrical. This results in an instantaneous temporary dipole. This dipole induces dipoles in neighboring molecules and leads to an attraction between the opposite charges in the dipoles

28
Q

Factors that affect the van der Walls’ force

A
  • Number of electrons - the more the stronger
  • The weaker the contact area, the stronger the induced dipole
  • Unbranched molecules have stronger van der Waals’ forces
29
Q

Trends in group VII

A

State progressively changes from gas to solid down the group
Reactivity decreases down the group
Mp/Bp increases down the group

30
Q

Electronegativity

A

The ability of an atom to attract electrons to itself in a bond

31
Q

Electronegativity’s of some elements

A 
F=4
O,N=3.5
Cl=3
Mg=1.2
Na=0.9
Al=1.5
32
Q

Permanent dipole

A

A small charge difference across a bong resulting from a difference in electro-negativities of the bonded atoms
Permanent dipole-dipole attractions occur in addition to van der Waals’ forces

33
Q

Trend in electronegativity

A

Electronegativity increases across a period and decreases down a group

34
Q

Use of electronegativity

A

Electronegativity can be used to predict bonding type
Electronegativity > 2.0 usually ionic
Electronegativity 1.0~2.0 usually polar
Electronegativity < 1.0 usually covalent

35
Q

Relative strengths of bonds and intermolecular forces

A

Ionic/covalent bonds 1000
Hydrogen bonds 50
Permanent dipole-dipole attraction 10
Van der Waals’ 1

36
Q

Polar covalent bond

A

A covalent bond with a permanent dipole

37
Q

Polar molecule

A

A molecule with an overall dipole when you take into account any dipoles across the bond

38
Q

When drawing hydrogen bonded molecules

A

Draw the hydrogen bond with a striped line
Draw the delta + and -
Position atoms either side of the H+ as far away from each other
Draw any lone pairs

39
Q

Hydrogen bonds

A

The attraction between a hydrogen atom on one molecule and a lone pair of another O, N or F molecule

40
Q

The reason water is not like the other Group VI hydrides

A

All group VI hydrides have a van der Waals’ attraction. On top of this, water has hydrogen bonding so its mp/bp are different

41
Q

The reason ice floats on water

A

In a solid, water forms a rigid structure held apart by hydrogen bonds. When this structure melts, the rigid hydrogen bonds break and the molecules move closer together

Chemistry (8 decks)

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