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AS Chemistry > 2.2.2 - Bonding and Structure > Flashcards

2.2.2 - Bonding and Structure Flashcards

1
Q

What does ionic bonding happen between?

A

A metal and a non metal.

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2
Q

Which way are electrons transferred?

A

From the metal atom to the non-metal atom.

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3
Q

What holds ionic compounds together?

A

Oppositely charged ions are formed, which are bonded together by strong electrostatic attraction.

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4
Q

In giant ionic lattices what is each ion surrounded by?

A

Oppositely charged ions.

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5
Q

How is a three dimensional giant ionic lattice formed?

A

Ions attract each other from all directions.

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6
Q

Can ionic compounds conduct electricity when solid?

A

No

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7
Q

Why can ionic compounds conduct electricity when in liquid state?

A

The ions are free to move around and conduct electricity.

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8
Q

What is ionic bonding?

A

electrostaic atracion between positive and negatively charged ions.

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9
Q

Why don’t ionic compounds conduct when solid?

A

Ions are in fixed position and cannot move to carry a charge.

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10
Q

Why can ionic compounds conduct when molten or in solution?

A

Ions are mobile and are therefore able to carry a charge.

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11
Q

Why do ionic compounds have high melting and boiling points?

A

Due to the strong electrostatic force of attraction between positively and negatively charged ions.

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12
Q

Why does MgO have require more energy to break down than NaCl?

A

Mg2+ and O2- bonds are stronger than Na- and Cl- bonds.

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13
Q

What do ionic lattices dissolve in?

A

Polar solvents, such as water.

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14
Q

What do polar bonds occur between?

A

Atoms that do not share electrons equally between them.

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15
Q

What is the result of a polar bond?

A

Results in the atom having very small charges on them.

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16
Q

What do polar water molecules do?

A

Attract the charged ions in the giant ionic lattice.

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17
Q

What is covalent bonding?

A

The strong electrostatic attraction between a shared pair of electrons and the nuclei of the bonded atoms.

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18
Q

What is a dative covalent bond?

A

A covalent bond in which the shared pair of electrons has been provided by one of the bonding atoms only.

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19
Q

What is an example of a substance which contains a dative covalent bond?

A

Ammonium ion. NH4+

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20
Q

What does the average bond enthalpy measure?

A

The energy required to break a covalent bond.

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21
Q

What does a greater average bond enthalpy mean?

A

The stronger the bond is, therefore more energy is required tot break it.

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22
Q

In shapes of molecules what angles are the biggest?

A

Lone pair lone pair angles

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23
Q

In shapes of molecules what angles are the smallest?

A

Bond pair - bond pair angles.

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24
Q

According to electron pair repulsion theory what happens to the bond angle when a lone pair is added?

A

The distances decreases by 2.5 degrees as the lone pair repels the bond pair.

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25
Q

In a linear molecule what is the bond angle?

A

180

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26
Q

If there are three bond pairs and no lone pairs what shape the molecule and what is its bond angle?

A

Trigonal planar and 120

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27
Q

What is the shape of a molecule and bond angle which has four bond pairs and no lone pairs?

A

Tetrahedral 109.5

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28
Q

What is the shape and bond angle of a molecule which has 3 bond pairs and 1 lone pair

A

Trigonal pyramidal 107

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29
Q

What is the shape and bond angle of a molecule which has two bond pair and two lone pairs?

A

Non linear 104.5

30
Q

What is the shape and bond angle of water?

A

Non-linear 104.5

31
Q

What is the shape and bond angle of a molecule with 5 electron pairs bonded around the central atom?

A

Trigonal bipyramidal 90

32
Q

What is the shape and bond angle of a molecule with 6 electron pairs around the central atom?

A

Octahedral 90

33
Q

What is electronegativity?

A

The ability of an atom to attract the bonding electrons in a covalent bond.

34
Q

What is the most electronegative element?

A

Fluorine

35
Q

How does the Pauling scale work?

A

The higher the value the more electronegative the element is.

36
Q

What makes a bond polar in covalent bonding?

A

The bonding electrons are pulled towards the more electronegative atom.

37
Q

What is a dipole?

A

The difference in charge between the two atoms caused by a shift in electron density in the bond.

38
Q

How is a bond more polar?

A

The greater the difference in electronegativity.

39
Q

Why are diatomic gases non polar?

A

Because both atoms have equal negativity and so electrons are attracted equally to both nuclei.

40
Q

Why are carbon and hydrogen ann-polar?

A

They have pretty similar electronegativity.

41
Q

What do polar bonds have?

A

Permanent dipoles.

42
Q

What does the arrangement of polar bonds in a molecule determine?

A

Whether or not the molecule will have an overall dipole.

43
Q

In what instance do polar molecules cancel out?

A

When the polar bonds are symmetrical - overall no dipole so non-polar.

44
Q

Example of cancelling out dipoles?

A

CO2

45
Q

In what case does the molecule have an overall dipole?

A

If the polar bonds are arranged unevenly across the whole molecule.

46
Q

Example of overall polar molecule?

A

H2O - negative charge is positioned more towards the oxygen atom.

47
Q

When is a substance purely covalent?

A

Diatomic gases because the electronegativity difference is 0 and so the bonding electrons are arranged completely evenly within the bond.

48
Q

How can you use electronegativity to predict the type of bonding?

A

The higher the difference in electronegativity, the more ionic in character the bonding becomes.

49
Q

What are intermolecular forces?

A

Forces between molecules.

50
Q

What is the strength of intermolecular forces?

A

They’re much weaker than covalent, ionic or metallic bonds.

51
Q

What are the three types of intermolecular forces?

A
  • London forces.
  • Permanet dipole-dipole interactions.
  • Hydrogen bonding
52
Q

What is the strongest intermolecular force?

A

Hydrogen bonding

53
Q

How is a temporary dipole formed in a charge cloud?

A

These clouds are moving really quick. At any particular moment, the electrons in an atom are likely to be more to one side than the other. This causes a temporary dipole.

54
Q

What does a temporary dipole cause?

A

Another temporary dipole in the opposite direction on a neighboring atom.

55
Q

How is the overall affect that the atoms are attracted to one another?

A

Dipoles are being constantly created and destroyed all the time as electrons are constantly moving.

56
Q

What do stronger London forces mean?

A

Higher boiling points.

57
Q

Why do larger molecules have stronger London forces?

A

They have a larger electron cloud.

58
Q

What do molecules with a greater surface area have?

A

A stronger London force because they have a bigger exposed electron cloud.

59
Q

What do molecules with stronger London forces have and why?

A

Higher boiling points as more energy is required to break the intermolecular forces.

60
Q

What kind of intermolecular forces do polar molecules have?

A

Permanet dipole-dipole interactions.

61
Q

When does hydrogen bonding occur?

A

When hydrogen is covalently bonded to fluorine, nitrogen or oxygen.

62
Q

What do molecules have which have hydrogen bonding?

A

-OH or -NH groups.

63
Q

How does hydrogen bonding affect the properties of a substance?

A

They are soluble in water and their density and viscosity are also affected.

64
Q

How is ice less dense than water?

A

Ice has more hydrogen bonds than liquid water and since the hydrogen bonds are relatively long, this makes the ice less dense than water.

65
Q

What is the main factor that determines the boiling point of a substance?

A

The strength of the induced dipole-dipole unless it has hydrogen bonds.

66
Q

What else increases the strength of the induced dipole-dipole?

A

Number of electrons in the molecule.

67
Q

Why are polar molecules soluble in water?

A

Water is a polar molecule, so only tends to dissolve other polar substances. Compounds with hydrogen bonds can from hydrogen bonds with water molecules so will be soluble.

68
Q

Why don’t simple covalent compounds conduct electricity?

A

Covalent substances are uncharged.

69
Q

Anomalous properties of water?

A
  • Less dense when ice as there is an open lattice due to longer hydrogen bonds.
  • Higher melting/boiling points than expected as hydrogen bonds require a larger amount of energy to overcome.
70
Q

Metallic bonding?

A

The strong electrostatic attraction between positive ions and negatively charged delocalised electrons in a regular lattice.

AS Chemistry (24 decks)

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