Brainscape's Knowledge GenomeTM


Top Flashcards (Certified)
All Flashcards

A Level Chemistry Edexcel > Topic 2: Bonding & Structure 1️⃣&2️⃣ > Flashcards

Topic 2: Bonding & Structure 1️⃣&2️⃣ Flashcards

1
Q

What are the typical physical properties of a metal

A
  • High melting temps
  • Good electrical conductivity
  • Good thermal conductivity
  • Malleability
  • Ductility
How well did you know this?
1
Not at all
2
3
4
5
Perfectly
2
Q

Define metallic bonding

A

The electrostatic force of attraction between the nuclei of metal cations and delocalised electrons

How well did you know this?
1
Not at all
2
3
4
5
Perfectly
3
Q

Which factors affects high melting temperature in metals

A

The number of delocalised electrons per cation plays a role in determining the melting temp : Group 1 have lower melting points than group 2
Size of cation also affects melting temp : Smaller the cation the smaller the atomic radii = increasing in melting temp

How well did you know this?
1
Not at all
2
3
4
5
Perfectly
4
Q

Explain electrical conductivity in metals

A

When a potential difference is applied across the ends of a metal the delocalised electrons will be attracted to and move towards the positive terminal of the cell - movement of electrical charge constitutes as an electric current

How well did you know this?
1
Not at all
2
3
4
5
Perfectly
5
Q

Explain thermal conductivity of metals

A

2 factors:
Delocalised electrons are free moving - can pass KE along the metal
Cations are closely packed - can pass KE from one cation to another

How well did you know this?
1
Not at all
2
3
4
5
Perfectly
6
Q

Explain malleability and ductility of metals

A

When stress is applied to a metal the layers of cations may slide over one another
Since the delocalised are free moving, they move with the cations and prevent strong forces of repulsion forming between cations in different layers

How well did you know this?
1
Not at all
2
3
4
5
Perfectly
7
Q

What is ionic bonding confined to

A

Solid materials consisting of a regular array of oppositely charged ions extending throughout a giant lattice network

How well did you know this?
1
Not at all
2
3
4
5
Perfectly
8
Q

What factors affect the strength of ionic bonding ?

A

Size of the ions : smaller= stronger
Charge on ions : larger = stronger

How well did you know this?
1
Not at all
2
3
4
5
Perfectly
9
Q

What are the physical properties of ionic compounds

A

High melting temperatures
Brittleness
Poor electrical conductivity when solid but good when molten
Often soluble in water

How well did you know this?
1
Not at all
2
3
4
5
Perfectly
10
Q

Explain the high melting temperatures of ioniccompounds

A

Many ions in a lattice and the combined electrostatic forces of attraction among all of the ions is large so a large amount of energy is required to overcome those forces sufficiently for the ions to break free from the lattice and be able to slide past eachother

How well did you know this?
1
Not at all
2
3
4
5
Perfectly
11
Q

Explain the brittleness of ionic compounds

A

If stress is applied to a crystal if an ionic solid then the layers of ions may slide over eachother
Ions of the same charge are now side by side and repel one another so the crystals break apart

How well did you know this?
1
Not at all
2
3
4
5
Perfectly
12
Q

Explain the electrical conductivity of ionic compounds

A

No delocalised electrons & ions are not free to move under the influence of an applied potential difference
Molten ionic compounds will conduct since the ions are now mobile and will migrate to the electrodes of the opposite charge

How well did you know this?
1
Not at all
2
3
4
5
Perfectly
13
Q

Explain the solubility of ionic compounds

A
  • Energy required to break apart the lattice structure and separate the ions can in some instances be supplied by the hydration of the separated ions produced.
  • Both positive and negative ions are attracted to water molecules because of the polarity that water molecules possess.
How well did you know this?
1
Not at all
2
3
4
5
Perfectly
14
Q

What is the primary evidence for the existence of ions

A

Ability if an ionic compound to conduct electricity and undergo electrolysis when either molten or in aqueous solution
- positive ions will migrate towards the negative electrode where they gain electrons and become atoms
- negative ions will migrate to wards the positive electrode where they lose electrons and become molecules

How well did you know this?
1
Not at all
2
3
4
5
Perfectly
15
Q

Define : covalent bond

A

The strong electrostatic attraction between the nuclei of two atoms and the bonding pair of electrons

How well did you know this?
1
Not at all
2
3
4
5
Perfectly
16
Q

What does an end-on overlap of two s-orbitals form ?

A

A sigma bond

How well did you know this?
1
Not at all
2
3
4
5
Perfectly
17
Q

What does an end-on overlap of 2 p-orbitals form ?

A

A sigma bond

How well did you know this?
1
Not at all
2
3
4
5
Perfectly
18
Q

What does a sideways overlap of 2 p-orbitals form ?
What is special about them?

A

A pi bond
- can’t be formed until an sigma bond has been formed which means pi bonds only exist between atoms that are joined by double/triple bonding

How well did you know this?
1
Not at all
2
3
4
5
Perfectly
19
Q

What is the reason for increased reactivity of alkenes compared to alkanes ?

A

The pi bond in ethene is weaker than the sigma bond

How well did you know this?
1
Not at all
2
3
4
5
Perfectly
20
Q

What is bond length

A

The distance between nuclei of the two atoms that are covalently bonded together

How well did you know this?
1
Not at all
2
3
4
5
Perfectly
21
Q

What is the relationship between bond length and bond strength

A

The shorter the bond length the greater the bond strength. This is a result of an increase in electrostatic attraction between the 2 nuclei and the electrons in the overlapping atomic orbitals

How well did you know this?
1
Not at all
2
3
4
5
Perfectly
22
Q

Define electronegativity

A

The ability of an atom to attract a bonding pair of electrons

How well did you know this?
1
Not at all
2
3
4
5
Perfectly
23
Q

What is the general trend in electronegativity

A

Decreased down the periodic table
Increases from left to right along the periodic table

How well did you know this?
1
Not at all
2
3
4
5
Perfectly
24
Q

How does electronegativity affect the distribution of electron density

A
  • If two atoms of the same element are bonded together by the overlap of atomic orbitals, the distribution of electron density will be symmetrical because their electronegativities are identical
  • However if the two bonded atoms have different electronegativities the distribution won’t be symmetrical - the end of the molecule with the higher electron density will have a slightly negative charge
How well did you know this?
1
Not at all
2
3
4
5
Perfectly
25
Define polar covalent bond
A type of covalent bond between two atoms where the bonding electrons are unequally distributed. Because of this, one atoms carries a slight negative charge and the other, a slight positive charge
26
What is a discrete (simple) molecule
An electrically neutral group of two or more atoms held together by chemical bonds
27
How and when is a **dative** covalent bond formed?
Formed when an empty orbital of one atom, overlaps with an orbital containing a non-bonding (lone pair) of electrons from another atom Examples : H3O+, NH4+, Al2Cl6
28
How is a dative covalent bond often showed in displayed formula
By an arrow starting from the atom providing the electron pair and going towards the atom with the empty orbital
29
The electron pair repulsion theory states that…
- the **shape of a molecule** is caused by repulsion between the pairs of electrons, both bond pairs and lone pairs that surround the central atom. - the electrons arrange themselves around the central atom so that the repulsion is minimised - lone pair-lone pair repulsion> lone pair-bond pair> bond pair-bond pair
30
When determining the shape of a molecule, how should you approach molecules with multiple bonds?
Treat each multiple bond as if it contained only one pair of electrons
31
What is the shape and bond angle of a molecule with 2 bonded pairs ?
Linear 180 °
32
What is the shape and bond angle of a molecule containing 3 bonded pairs
Trigonal planar 120 °
33
What is the shape and bond angle of a molecule containing 4 bonded pairs ?
Tetrahedral 109.5 °
34
What is the shape and bond angle of a molecule containing 6 bonded pairs ?
Octahedral 90 °
35
What is the shape and bond angle of a molecule containing 2 bonded pairs and 1 lone pair ?
Non-linear 117.5 °
36
What is the shape and bond angles of a molecule containing 3 bonded pairs and 1 lone pair
Trigonal pyramidal 107 °
37
How do you calculate the bond angle for a molecule contains lone pair(s)?
Find the total number of electron groups (bonded pairs and lone pairs) Say its 2 bonded pairs and 2 lone pairs, that would be 4 electron groups so would take bond angle of a tetrahedral molecule (109.5 °) However for each lone pair the bond angle is reduced by ~2.5 ° 109.5- (2.5 x2) = 104.5 °
38
What is a dipole
The drift of bonded electrons towards the more electronegative element that results in a separation of charge
39
What does the overall dipole (whether a molecule is polar/non-polar) depend on ?
It’s shape and relative angles The individual dipoles can either reinforce one another or cancel each other out - if cancellation occurs, the molecule will have no overall dipole and is said to be ‘non-polar’ - if the dipoles reinforce eachother, the molecule will possess an overall dipole and is said to be ‘polar’
40
How is an instantaneous dipole created?
Electron density fluctuates overtime in a non polar molecule. If at any instance the electron density becomes **unsymmetrical** a dipole will be generated The end of the molecule with an increased electron density will have a partial negative charge (δ-) and the area where the electron density has decreased will have a partial positive charge (δ+)
41
How is the induced dipole created and how is that in turn responsible for the London forces?
The δ+ area on molecule A (with the instantaneous dipole) approaches another molecule (B) Electron density of molecule B is pulled to the left generating a partial negative charge on one side and a partial positive charge on the other Since it was the dipole of A that lead to the induction of dipole in B, the two molecules are arranged so they can interact with eachother ( interaction is responsible for London forces)
42
Describe & explain the trend in London forces as electrons per molecule increases
London forces increase as the number of electrons in a molecule increases why? The more electrons in a molecule the greater the fluctuation in electron density and the larger the instantaneous and induced dipoles created
43
Describe the trend in London forces as the shape and size of the molecule increases
The more points of interaction between molecules the greater the overall London force
44
How do molecules with permanent dipoles interact and why is this interaction often weaker than that between instantaneous-induced dipoles?
If the dipoles are aligned correctly then there will be a favourable interaction and the two molecules will attract one another Interaction is often weaker **due to random movement of molecules** the **dipoles are not always aligned to produce a favourable interaction**
45
Describe hydrogen bonding through nitrogen
All compounds containing an N-H group can form intermolecular hydrogen bonds Example : ammonia, NH3+
46
Describe hydrogen bonding through fluorine
The only fluorine compound with intermolecular hydrogen bonding is hydrogen fluoride
47
Define the hydrogen bond
An **intermolecular interaction** between a hydrogen atom of a molecule bonded to an atom which is more electronegative than hydrogen and another atom in the same or different molecule
48
What are the reasons for the increase in boiling temp with increasing molecular mass in unbranched alkanes?
As molecular mass increases the number of **electrons per molecule** increases and so the **instantaneous induced dipoles** increase As length of carbon chain increases the **number of points of contact** between adjacent molecules increases. Instantaneous dipole-induced dipole forces exist at each point of contact between molecules so the more points of contact the greater the overall intermolecular (London) force of attraction
49
How does branching affect the boiling points of alkanes
The more branching in a molecule, the fewest points of contact between adjacent molecules. Leads to a decrease in the overall intermolecular force of attraction between molecules and a decrease in boiling temp
50
What is the only significant intermolecular interaction in alkane molecules
London forces
51
What are the significant intermolecular interactions between alcohol molecules ?
Contain an OH group so can form intermolecular hydrogen bonds in addition to London forces
52
How does alcohols additional ability to form hydrogen bonds have an affect on the boiling temperature compared to alkanes
Hydrogen bonding provides an additional force of attraction that increases the energy required to separate the molecules
53
What are the anamolous properties of water
1. It has a relatively high melting and boiling temp for a molecule with so few electrons 2. The density of ice at 0°C is less of that of water at 0°C 
54
Explain the anomalous property of water : high melting and boiling temps
Hydrogen bonds in water molecules are relatively strong and as a result the intermolecular forces of attraction in water are greater than would be expected from the number of electrons
55
Why is the boiling point of water greater than that of HF even though the bond strength for HF is greater? (They both have same number of electrons)
HF forms an average of one hydrogen bond per molecule whereas water forms 2 so the hydrogen bonding is much more extensive in water Not all the hydrogen bonds in HF are broken on vaporisation since HF is substantially polymerised, even in the gas phase
56
Explain the anomalous property of water: the density of ice
The molecules in ice are arranged in rings of 6, held together by hydrogen bonds The structure creates a large areas of open space inside the rings so when ice melts the ring structure is destroyed and the average distance between molecules decreases, causing an increase in density
57
What are the conditions that must be met when choosing a suitable solvent ?
The solute particle must be separated form eachother and then become surrounded by solvent particles The forces of attraction between the solute and the solvent must be strong enough to overcome the solvent-solvent forces & the solute-solute forces
58
When dissolving ionic solids, how is the energy required to seperate the ions in the solid supplied?
By the hydration of the ions
59
Describe and explain the process of dissolving ionic solids in water
The δ– end of the water molecules attract the positively charged solid ions sufficiently to remove them from the lattice. They become surrounded by water molecules (this interaction is called an ion-dipole interaction) The δ+ end of the water molecules attract the negatively charged solid ions in the same way and they too become surrounded by water molecules Energy released is called the ‘hydration energy’
60
Which compounds can form hydrogen bonds with water
Alcohols However the solubility of alcohols in water decreases with increasing hydrocarbon chain length as London forces predominate between alcohol molecules
61
Which compounds cannot form hydrogen bonds with water
Non polar molecules such as alkanes Halogenoalkanes (much more soluble in ethanol) Some polar molecules
62
Why do **some** polar molecules have limited solubility in water
Either they don’t form hydrogen bonds with water or the hydrogen bonds they form are weak compared to those of water The forces of attraction between the polar molecule and water molecules are not large enough to replace the relatively strong hydrogen bonding between water molecules
63
Describe a metallic lattice
Metallic lattices are composed of a regular arrangement of positive metal ions surrounded by delocalised electrons
64
What are the four most common giant covalent lattices
Diamond Graphite Graphene Silicon (IV) oxide
65
Describe the bonding and properties of diamond
Each carbon atom forms four sigma bonds to four other carbon atoms Gaint 3D tetrahedral arrangement (bond angles all 109.5) Very hard and high melting points due to strong C-C bonding throughout
66
Describe the bonding of graphite
Each carbon is bonded to 3 others by sigma bonds forming **interlocking hexagonal rings** The fourth electron of each carbon atom is in the p-orbital The carbon atoms are close enough for the p-orbitals to overlap with eachother and produce a cloud of delocalised electrons above and below the plane of the rings
67
Describe the properties of graphite
Can be **used as a solid lubricant** since the layers slide easily over eachother - lubrciating properties are as a result of absorbed gases on the surface of the carbon atoms Fairly **good conductor of electricity** - delocalised electrons between the layers are free to move under the influence of applied potential difference (however can only conduct parallel to its layers) **High melting point** for same reason as diamond
68
Describe the structure and bonding of graphene
One atom thick (pure carbon) Carbon atoms are bonded in the same way as graphite
69
Describe and explain the physical properties of molecular solids
**Low melting and boiling temperatures** - in order to melt only necessary to overcome the intermolecular forces of attraction, no need to break covalent bonds - Little energy is needed to break down lattice as intermolecular forces of attraction are much weaker than covalent bonds **Generally insoluble** but may dissolve if hydrogen bonding is possible **Non- conductor**

A Level Chemistry Edexcel (20 decks)

Brainscape helps you reach your goals faster, through stronger study habits.
© 2025 Bold Learning Solutions. Terms and Conditions