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AS CHEMISTRY > BONDING > Flashcards

BONDING Flashcards

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1
Q

Define ionic bonding

A

The electrostatic force of attraction

between oppositely charged ions

formed by electron transfer

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2
Q

Metal atoms ___ electron to form _ve ions

Non-metal atoms ___ electron to form _ve ions

A

Metal lose to form +ve

Non-metal gain to form -ve

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3
Q

How do smaller ions impact melting points

A

High melting points

In smaller ions

Because ionic bonding is stronger

Or a higher charge

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4
Q

Define covalent bonding

A

A shared pair of electrons

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5
Q

Define dative covalent bonding / Coordinate bindi g

A

A shared pair of electrons

Donated from only one of the bonding atoms

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6
Q

Give the 3 common dative covalent molecules

A

NH4 +

H3O +

NH3BF3

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7
Q

How is a dative covalent bond represented

A

An arrow

From the atom providing the lone pair

To the deficient atom

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8
Q

Define metallic bonding

A

Electron static force of attraction

Between positive metal ions

And delocalised electrons

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9
Q

List the 3 factors effecting the strength of a metallic bond

A

Proton number

Delocalised electron number (per atom)

Size of ion

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10
Q

How does proton number affect the strength of metallic bonding

A

More protons

Stronger bond

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11
Q

How does delocalised electron number affect the strength of metallic bonding

A

More delocalised electrons

Stronger bond

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12
Q

How does ion size affect the strength of metallic bonding

A

Smaller ion

Stronger bond

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13
Q

Give an example of a giant ionic lattice molecule

A

Sodium chloride

Magnesium oxide

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14
Q

Give an example of a simple molecular molecule

A

Iodine

Ice

Carbon dioxide

Water

Methane

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15
Q

Give an example of a macromolecular molecule

A

Diamond

Graphite

Silicon dioxide

Silicon

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16
Q

Give an example of a giant metallic lattice molecule

A

Magnesium

Sodium

All metals

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17
Q

Outline the properties of giant ionic lattices

A

High melting and boiling point

Soluble in water

Poor conductor when solid

Good conductor when molten/dissolved

Crystalline solid

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18
Q

Outline the properties of a simple molecular molecule

A

Low melting and boiling point

Poor solubility in water

Poor conductor when solid and molten

Mostly gas or liquid

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19
Q

Outline the properties of a macromolecule

A

High melting and boiling point

Insoluble in water

Diamond/Sand are poor conductors when solid but graphite is good

Poor conductors when molten

Solids

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20
Q

Describe the properties of a metallic molecule

A

High melting and boiling point

Insoluble in water

Good conductors when solid and molten

Shiny

Malleable

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21
Q

Outline a linear molecule. Give an example.

A

2 bonding pairs

No lone pairs

180 bond angle

CO2, CS2, HCN, BeF2

22
Q

Outline a trigonal planar molecule. Give an example.

A

3 bonding pairs

No lone pairs

120 bond angle

BF3, AlCl3, SO3, NO3 -, CO3 2-

23
Q

Outline a tetrahedral molecule. Give an example.

A

4 bonding pairs

No lone pairs

109.5 bond angle

SiCl4, SO4 2-, ClO4 -, NH4 -

24
Q

Outline a trigonal pyramidal molecule. Give an example.

A

3 bonding pairs

1 lone pair

107 bond angle

NCl3, PF3, ClO3, H3O +

25
Q

Outline a bent molecule. Give an example.

A

2 bonding pairs

2 lone pairs

104.5 bond angle

OCl2, H2S, OF2, SCl2

26
Q

Outline a trigonal bipyramidal molecule. Give an example.

A

5 bonding pairs

No lone pairs

120 and 90 bond angles

PCl5

27
Q

Outline a octrahedral molecule. Give an example.

A

6 bonding pairs

No lone pairs

90 bonding angle

SF6

28
Q

Outline the structure for explaining the shape of a molecule

A
  1. state number of bonding pairs and lone pairs
  2. State electron pairs repel to a position of minimum repulsion
    (3. State lone pairs repel more than bonding pairs)
  3. State actual shape and bonding angle/s
29
Q

How much do lone pairs reduce bond angles

A

2.5

30
Q

Outline a square planar molecule. Give an example.

A

4 bonding pairs

2 lone pairs

Bond angle 90

XeF4

31
Q

Outline the shape of BrF5

A

Bond angle 89

4 bond pairs

2 lone pairs

32
Q

Outline the shape of I3 -

A

Bond angle 180

2 bonding pairs

3 lone pairs

33
Q

Outline the shape of ClF3

A

Bond angle 89

3 bonding pairs

2 lone pairs

34
Q

Outline the shape of SF4 of IF4 +

A

Bond angles of 119 and 89

3 bonding pairs

1 lone pair

35
Q

Define electronegativity

A

The relative tendency

of an atom in a covalent bond

in a molecule to attract electrons

in a covalent bond to itself

36
Q

List the 4 most electronegative elements

A

Fluorine

Oxygen

Nitrogen

Chlorine

37
Q

What factors affect electronegativity

A

Proton number

Atomic radius

Shielding

38
Q

Outline how electronegativity changes across the period

A

Increases

Larger proton numbers

Smaller atomic radius

Electrons closer to nucleus

39
Q

Outline how electronegativity changes down a group

A

Decreases

More shells

More shielding

Electrons further from the nucleus

40
Q

How does electronegativity effect the type of bonding in a compound

A

Similar electronegativity
Covalent bond

Large difference in electronegativity
Ionic bond

41
Q

Outline how a permanent dipole (polar covalent) bond forms

A

Elements in the bond have different electronegativity

Difference of around 0.3 to 1.7

Causing an unequal distribution in electrons

And thus a charge separation

(Delta) + and (delta) -

42
Q

Outline the effects of symmetry in molecules

A

All bonds are identical
With no lone pairs

Cannot be a polar molecule
As the any dipoles cancel out
Due to symmetry

43
Q

When does van der waals occur

A

Between all molecular substances

And Nobel gases

Not in ionic substances

44
Q

Describe how van der waals (induced dipole) occurs

A

Electron density fluctuates

Parts of the molecule become more/less negative

Induces neighbouring molecules to do the same (induced dipole)

The induced dipole is the opposite to the original one

45
Q

Outline the factors effecting van der waals

A

More electrons
More likely to form

Increased van der waals
Higher boiling points

Long chain alkanes more likely
More surface area

Eg halogens increase down group due to more van der waals

46
Q

When does permanent dipole-dipole occur

A

Between polar molecules

47
Q

Outline permanent dipole dipole forces

A

Stronger than van der waals

Further increase boiling point

Molecules are asymmetrical
and have significant electronegativity difference between atoms

48
Q

When does hydrogen bonding occur?

A

Between hydrogen atoms and either:
Oxygen
Nitrogen
fluorine

With a lone pair available

49
Q

Outline hydrogen bonding

A

Strongest intermolecular bond

Causes very high boiling points (eg water is a liquid not gas)

Alcohols, carboxylic acids, proteins, amides all form hydrogen bonds

50
Q

Order the intermolecular forces from strongest to weakest

A

Hydrogen bonding

Permanent dipole dipole

Van der waals

AS CHEMISTRY (17 decks)

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