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Chemistry A Level Pearson's Edexcel > Bonding & Structure > Flashcards

Bonding & Structure Flashcards

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1
Q

Define ionic bonding

A

Electrostatic attraction between oppositely charged ions

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2
Q

Define metallic bonding

A

Electrostatic attraction between positive metal cations and delocalised electrons

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3
Q

Define covalent bonding

A

Electrostatic attraction between protons in two nuclei and a shaired pair of electrons between them

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4
Q

Explain why water has a low boiling point.

A

Weak forces between the H2O molecules (weak “intermolecular forces”)
Little energy is needed to overcome them.

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5
Q

What is required for a substance to conduct electricity? Using this, explain why water cannot conduct electricity.

A

For a substance to conduct electricity, it must have freely moving charges. Although the H2O molecules can move, they are neutral.

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6
Q

What does covalent bond strength depend on?

A

Bond length - shorter bonds, stronger bonds
Bond order (more e- pairs, stronger attraction)

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7
Q

Define allotropes.

A

Forms of an element with different structures.

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8
Q

What are the 4 main carbon allotropes?

A

Diamond
Graphite
Graphene
Nanotubes & fullerenes

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9
Q

State the properties of diamond

A

V. high MP
Hard - good for cutting
Electrical insulator
Great conductor of sound/heat

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10
Q

State the properties of graphite

A

High MP
Soft
Electrical conductor

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11
Q

What does VSEPR stand for?

A

Valence-Shell Electron-Pair Repulsion

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12
Q

What is the shape and bond angle around a molecule with two bond pairs?

A

Linear, 180*

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13
Q

What is the shape and angle around an atom with three bond pairs?

A

Trigonal planar, 120*

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14
Q

What is the shape and angle around an atom with four bond pairs?

A

Tetrahedral, 109.5

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15
Q

What is the shape and angle around an atom with five bond pairs?

A

Trigonal bipyramid, 120*

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16
Q

What is the shape and angle around an atom with six bond pairs?

A

Octahedral, 90*

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17
Q

Explain why specific numbers of bond/electron pairs result in their respective angles & shapes.

A

Electron pairs separate as far as possible
To minimise electron pair repulsion

18
Q

What is the shape and angle around an atom with three bond pairs and one lone pair?

A

Trigonal pyramid, 107*

19
Q

What is the shape and angle around an atom with two bond pairs and two lone pairs?

A

V-shaped / Bent, 104-5*

20
Q

Why do the angles change when lone electron pairs are introduced?

A

Lone electron pairs repel bonding electron pairs more than bonding electron pairs repel bonding electron pairs.

21
Q

Why are some bonds polar?

A

A difference in electronegativity between atoms in a bond result in a polar bond.

22
Q

How can you tell experimentally if a molecule is polar?

A

Deflecting jet test:
polar liquids will deflect towards a charged rod

23
Q

What are the limitations to a delfecting jet test?

A

Doesn’t work with (s), (g) and (aq)
but… could dissolve into a non-polar solvent.

24
Q

What does the strength of london forces depend on?

A

Number of electrons per molecule - more e-, stronger attraction
Shape of molecule - closer packing, better attraction
Extent of delocalisation

25
Q

What are permanenet dipole-dipole forces?

A

Permanent dipole-dipole forces are the weak intermolecular forces of attraction that arise between permanently polar molecules. These forces are between the delta positive end of one polar bond with the delta negative end of another polar bond.

26
Q

What is hydrogen bonding?

A

Hydrogen bonding is electrostatic force of attraction between a hydrogen atom which is covalently bound to a more electronegative “donor” atom or group, and another electronegative atom bearing a lone pair of electrons—the hydrogen bond acceptor

27
Q

Why can hydrogen bonding only occur with certain elements?

A

Only occurs with nitrogen, oxygen or fluorine as the atom needs to be electronegative enough to make the X-H bond polar enough to expose the H nucleus.

28
Q

List different IMFs in order of strength (strongest to weakest)

A

Hydrogen bonding
Dipole-dipoles
London forces

29
Q

Define miscibility.

A

Miscibility is the property of two substances to mix in all proportions, forming a homogeneous mixture.

30
Q

How can you tell if two substances are miscible?

A

A substance is miscible if the bonds formed are equal to or greater than, in terms of strength, than the bonds broken.

31
Q

What affects the strength of ionic bonds and how?

A

Ionic radius - increasing radius decreases strength
Ionic change - increasing change increases attraction

32
Q

What is formed when an atom losses or gains an electron? What process does this?

A

Ions are formed
Cations from losing electrons in oxidation
Anions from gaining electrons in reduction

33
Q

Describe the general properties of ionic compounds.

A

High melting and boiling points
Soluble in polar solvents
Electrically conductive when molten or dissolved in water

34
Q

Define electronegativity.

A

Tendency of an atom to attract bonding electrons

35
Q

Describe the different types of intermoleuclar forces and what affects the strength of them.

A

London forces (Instantaneous dp - induced dp) - greater electron density, increases strength
Permanent dipoles - strength increases with increasing polarity in the molecule
Hydrogen bonds - increasing charge increases bond strength

36
Q

How does hydrogen bonding arise in molecules such as H20, NH3 or liquid HF?

A

When bonded to hydrogen, oxygen and nitrogen and flourine have a large enough difference in electronegativity such that a polar bond is formed. This allows the hydrogen to be positively charged enough to form hydrogen bonds to lone pairs on other atoms.

37
Q

How does the strength of hydrogen bonds compared to other IMFs?What properties does this effect and how?

A

Hydrogen bonds are far stronger than other IMFs and therefore take more energy to break. This results in an increase in melting/boiling points for substances that contain hydrogen bonds.

38
Q

How do hydrogen bonds affect the density of water/ice?

A

When freezing ice, more hydrogen bonds form between molecules which causes the molecules to space out and therefore decrease the density of water.

39
Q

How does bonding vary with different shapes of chains and chain length in alkanes? How does this affect melting/boiling points?

A

As chain length increase, london forces become stronger therefore increasing melting/boiling points.
Straighter thinner chains have stronger london forces compared to branched chains due to their ability to pack closer together.

40
Q

How does the boiling point vary between alcohols and alkenes with similar numbers of electrons? (e.g. methanol and ethane (both 18 e-))

A

Alcohols are polar and are able to form hydrogen bonds
Both alcohols and alkenes have london forces
DPDP and hydrogen bonds are much stronger therefore alcohols have higher boiling points

41
Q

Describe and explain the trend in boiling points between different hydrogen halides.

A

As the size of the halogen increases, the greater the boiling point due to the increase in london forces.

Chemistry A Level Pearson's Edexcel (17 decks)

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