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CHEMISTRY > >>3.1 - Bonding✔ > Flashcards

>>3.1 - Bonding✔ Flashcards

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1
Q

When can an induced dipole form?

A
  • Can form when the electron orbitals around a molecule are influenced by another charged particle.
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2
Q

What does the shape of a molecule depend on?

A
  • Determined by the number of electron pairs around the central atom.
  • Each electron pair naturally repels each other so that the largest bond angle possible exists between the covalent bonds.
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3
Q

What affect do lone pairs have around the central atom?

A
  • Provide additional repulsive forces which changes the bond angle.
  • For every lone pair present the bond angle between covalent bonds is reduced by 2.5°.
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4
Q

What 3 things should be considered for determining the shape of a molecule?

A
  • Find the number of electron pairs.
  • Determine how many of the pairs are bonding pairs and how many are lone pairs.
  • Bonding pairs indicate the basic shape and lone pairs indicate any additional repulsion.
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5
Q

Describe how to calculate the number of lone pairs for a molecule
(Use CH₄ as an example)

A

⓵Find the group number of the central atom.
⓶Add the group number to the number of bonded pairs.
⓷If charge is + then subtract 1, if charge is - then add 1 and if no charge then don’t add or subtract anything.
⓸Divide this number by 2 to get total electron pairs.
⓹Do electronpairs - bondedpairs to get lone pairs.

4+4=8
8÷2=4
4-4= 0LP,4BP

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6
Q

What is the shape, bond angle and diagram for a molecule containing 4 electron pairs, all 4 of which are bonded?
e.g: CH₄

A
  • Tetrahedral (symmetrical.)
  • 109.5°
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7
Q

What is the shape, bond angle and diagram for a molecule containing 4 electron pairs, 3 of which are bonded pairs and 1 of which is a lone pair?
e.g: NH₃

A
  • Pyramidal (asymmetrical.)
  • 107°
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8
Q

What is the shape, bond angle and diagram for a molecule containing 4 electron pairs, 2 of which are bonded pairs and 2 of which are lone pairs?
e.g: H₂O

A
  • V-Shape (asymmetrical.)
  • 104.5°
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9
Q

What is the shape, bond angle and diagram for a molecule containing 2 electron pairs, both of which are bonded pairs?
e.g: BeCl₂

A
  • Linear (symmetrical.)
  • 180°
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10
Q

What is the shape, bond angle and diagram for a molecule containing 3 electron pairs, all 3 of which are bonded pairs?
e.g: BCl₃

A
  • Trigonal planar (symmetrical.)
  • 120°
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11
Q

What is the shape, bond angle and diagram for a molecule containing 5 electron pairs, 3 of which are bonded pairs and 2 of which are lone pairs?
e.g: ClF₃

A
  • T-Shaped.
  • 86°
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12
Q

What is the shape, bond angle and diagram for a molecule containing 5 electron pairs, 2 of which are bonded pairs and 3 of which are lone pairs?

e.g: I₃⁻

A
  • Linear.
  • 180°
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13
Q

What is the shape, bond angle and diagram for a molecule containing 6 electron pairs, all 6 of which are bonded pairs?
e.g: SF₆

A
  • Octahedral (symmetrical.)
    *90°
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14
Q

What is the shape, bond angle and diagram for a molecule containing 6 electron pairs, 4 of which are bonded pairs and 2 of which are lone pairs?

e.g: ICl₄⁻

A
  • Square Planar.
  • 90°
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15
Q

What is electron pair repulsion theory?

A
  • Electrons always repel as far as possible.
  • Lonepair/Lonepair repulsion is greater than Lonepair/Bondpair repulsion which is greater than Bondpair/Bondpair repulsion.
  • When there are no lone pairs, the bonding pairs in a molecule repel equally.
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16
Q

Covalent bonds are directional in space, what does this cause?

A
  • Hence molecules have a set shape depending on the number of electron pairs around the central atom.
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17
Q

What is the shape, bond angle and diagram for a molecule containing 5 electron pairs, all 5 of which are bonded pairs?
e.g: PF₅

A
  • Trigonal Bypyramidal.
  • 120° and 90°
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18
Q

What is the shape, bond angle and diagram for a molecule containing 5 electron pairs, 4 of which are bonded pairs and 1 of which is a lone pair?
e.g: SF₄

A
  • Modified trigonal Bypyramidal.
  • 118° and 89°
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19
Q

What can polar molecules with a permanent dipole do?

A
  • Can align to form a lattice of molecules similar to an ionic lattice.
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20
Q

What is the structure and bonding in ice?

A
  • Open lattice structure.
  • Covalent bonding between atoms but hydrogen bonds between molecules.
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21
Q

Describe in detail ionic bonding:

A
  • Occurs between a metal and a non-metal.
  • Electrons are transferred from the metal to the non-metal to achieve full outer shells.
  • When the electrons are transferred it creates charged particles called ions. These oppositely charged ions attract through electrostatic forces to form a giant ionic lattice.
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22
Q

How is sodium chloride formed?

A
  • Ionic compound formed from Na+ and Cl- ions.
  • Sodium loses an electron and chlorine gains an electron to produce ions with a full outer electron shell.
  • These then attract into an ionic lattice.
    *Electrostatic attraction between oppositely charged ions in a lattice.
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23
Q

What are the formulas for the following compounds ions?
≫Sulfate:
≫Hydroxide:
≫Nitrate:
≫Carbonate:
≫Ammonium:

A
  • SO₄²⁻
  • OH⁻
  • NO₃⁻
  • CO₃²⁻
  • NH₄⁺
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24
Q

Describe covalent bonding:

A
  • Form between two non-metals.
  • Electrons are shared between the two outer shells in order to achieve a full outer shell.
  • There are electrostatic attractions between the shared pair of electrons and the positive nucleus.
    *Multiple bonds contain multiple pairs of electrons.
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25
Q

How can covalent bonding be shown?

A
  • Dot and cross diagrams.
  • Covalent bonds shown with a straight line.
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26
Q

Describe dative bonding:

A
  • Dative bonds form when both of the electrons in the shared pair are supplied from a single atom.
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27
Q

How can co-ordinate bonds be shown?

e.g: NH4+

A
  • Indicated using an arrow from the lone electron pair to the atom/molecule that its being donated too.
  • Ammonia has a lone pair of electrons that can form a dative bond with a H⁺ ion to produce NH₄⁺
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28
Q

What can be said about a dative bond once it has been formed?

A
  • It is treated as a standard covalent bond as it reacts in exactly the same way.
29
Q

Describe metallic bonding:

A
  • Consists of a lattice of positively charged ions surrounded by a ‘sea’ of delocalised electrons.
  • This produces a very strong electrostatic force of attraction between the oppositely charged particles.
30
Q

Metallic Bonding:
Describe the affect on force if…

≫Positive metal ion has a bigger charge:

≫Positive metal ions are larger:

A
  • The greater the charge on the ion, the stronger the electrostatic attraction as more electrons are released into the ‘sea’.
  • Ions that are larger in size produce a weaker electrostatic attraction due to their greater atomic radius.
31
Q

What are examples of physical properties and what do they depend on?

A
  • Boiling point, melting point, solubility and conductivity.
  • These depend on the type of bonding and the crystal structure of the compound.
32
Q

What are the four main types of crystal structure?

A
  • Giant Ionic.
  • Giant Metallic.
  • Simple Molecular.
  • Giant Macromolecular
33
Q

Describe the melting point and boiling points of substances with an ionic crystal structure and explain why:

A
  • High MP+BP.
  • Strong electrostatic attractions holding the ionic lattice together between the ions require large amounts of energy to overcome.
34
Q

Describe the electrical conductivity in substances with an ionic crystal structure and explain why:

A
  • When molten or aqueous they can as the ions aren’t held in a lattice and are free to move and carry a flow of charge.
  • In solid they cannot as ions are held in a lattice.
35
Q

Why are ionic substances often brittle?

A
  • When the layers of alternating charges are distorted, like charges repel breaking apart the lattice into fragments.
36
Q

Describe the electrical conductivity of substances with a metallic structure and explain why:

A
  • Good conductors.
  • The ‘sea’ of delocalised electrons are able to move and carry charge throughout the structure.
37
Q

Why are metals malleable?

A
  • Layers of positive ions are able to slide over one another.
  • The delocalised electrons prevents fragmentation as they can move around the lattice.
38
Q

Describe the melting points of metallic substances and explain why:

A
  • High melting points.
  • The electrostatic forces of attraction between the positive ions and delocalised electrons are very strong and therefore require a lot of energy to overcome.
39
Q

What do substances with a simple molecular structure consist of?

A
  • Consist of covalently bonded molecules held together with weak van Der Waals forces.
40
Q

Describe the melting points of structures that are simple molecular and explain why:

A
  • Low MP+BP.
  • Van Der Waals forces between Molecules are very weak and not much energy is required to overcome them.
41
Q

Water has a simple molecular structure, yet has an unusually high boiling point, why is this?

A
  • Due to the presence of hydrogen bonding.
42
Q

Describe the conductivity of simple molecular substances and explain why:

A
  • Very poor conductors.
  • Their structures contain no delocalised electrons thus cant carry charge throughout the structure.
43
Q

Describe the melting point of substances with a giant macromolecular structure and explain why:

A
  • Very high MP+BP.
  • Each atom has multiple covalent bonds which are very strong requiring a large amount of energy to break.
44
Q

Describe the structure of graphite:

A
  • Giant macromolecular structure.
  • Each carbon atom is bonded to 3 other carbon atoms.
  • Forms hexagonal layers.
45
Q

Describe and explain the two key properties of graphite:

A
  • Conducts electricity because each carbon atom bonds to only three others meaning there are delocalised electrons that can carry charge.
  • Soft because there are weak IMF between the hexagonal layers.
46
Q

Why can’t diamond conduct electricity?

A
  • Each carbon atom bonds to 4 other carbon atoms therefore there aren’t any delocalised electrons to carry charge through the structure.
47
Q

What is bond polarity?

A
  • Where the negative charge around a covalent bond is not evenly spread around the orbitals of the bonded atoms.
48
Q

What is electronegativity defined as?

A
  • The power of an atom to attract electrons towards itself within a covalent bond.
49
Q

What determines an atoms ‘power’ of electronegativity?

A
  • Its size (atomic radius.)
  • Nuclear Charge.
  • Electron Shielding.
50
Q

Describe the trend of electronegativity along a period and down a group:

A
  • Increases along a period as atomic radius decreases and nuclear charge increases for the same shielding.
  • Decreases down a group as shielding increases and atomic radius increases.
51
Q

What are the electronegativities of…
≫F:
≫Cl:
≫O:
≫N:
≫C:
≫H:

A
  • 4.0
  • 2.8
  • 3.5
  • 3.0
  • 2.5
  • 2.1
52
Q

When is a polar bond formed?

A
  • If the two atoms that are bonded have different electronegativities.
53
Q

What is a permanent dipole?

A
  • An unequal sharing of electrons in a covalent bond leading to 𝛿+ and 𝛿- regions.
  • This is due to a difference in electronegativities.
54
Q

Why is hydrogen fluoride a polar molecule?

A
  • Fluorine is a lot more electronegative than hydrogen so electrons are drawn towards fluorine.
55
Q

Why are molecules such as methane effectively non-polar?

A
  • They are symmetrical so each of the individual dipoles cancel each other out effectively leaving it non-polar.
56
Q

What are the three main types of intermolecular forces?

A
  • Van der Waals forces.
  • Permanent Dipole.
  • Hydrogen Bonding.
57
Q

Where do Van der Waals forces occur?

A
  • Exists between all molecules whether polar or non-polar and even noble gasses.
58
Q

How does a Van der Waals force form?

A
  • VDW form because of electron movement in the first molecule creating a temporary dipole.
  • This induces a dipole in another molecule.
  • This forms an attraction of 𝛿+ and 𝛿- in adjacent molecules.
59
Q

Describe and explain two factors affecting the strength of Van der Waals forces:

A
  • Larger the Mr of the molecule the more electrons present and therefore stronger the force between molecules.
  • Straight chain molecules experience stronger VDW forces than branched chain molecules as they can line up and pack closer together. This reduces the distance over which the force acts and so it is stronger.
60
Q

Where do permanent dipole-dipole forces occur?

A
  • This type of IMF acts between molecules with a polar bond. The 𝛿+ and 𝛿- regions attract each other and hold the molecules together in a lattice like structure.
61
Q

Sketch a diagram of the IMF between water molecules and ethanol:

A
62
Q

What do hydrogen bonds form between?

A
  • Only form between hydrogen and one of the three most electronegative atoms: N, O or F.
63
Q

How does a hydrogen bond form?

A
  • Due to uneven sharing of electrons because of the different electronegativities hydrogen becomes 𝛿+ and N, O or F becomes 𝛿-.
  • Lone pair of electrons from N, O or F forms a ‘bond’ with 𝛿+ H in another molecule.
64
Q

What can be said for molecules held together with hydrogen bonds compared to similar sized molecules without hydrogen bonding?

A
  • Molecules held together with hydrogen bonds have much higher MP+BP.
65
Q

What are the 3 key stages when drawing a diagram to show hydrogen bonding?

A

⓵Add polar charges to molecules.
⓶Add lone pairs of electrons.
⓷Show and label the hydrogen bond using a dotted line, the angles of the molecules must be in 180 degrees with the bond.

66
Q

Why does ice have a lower density than water?

A
  • When water freezes the hydrogen bonds hold the water molecules in a more open 3D crystalline structure, the water molecules are further apart from each other than they were in liquid therefore ice has a lower density.
  • In water the hydrogen bonds are constantly breaking and reforming.
67
Q

What is another name for a positive ion?

A
  • Cation
68
Q

What is another name for a negative ion?

A
  • Anion
69
Q

What does the word polyatomic mean?

A
  • Numerous atoms (E.g. SO42-)

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