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

Topic 2: Bonding & Structure Flashcards

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1
Q

What are the typical physical properties of a metal

A
  • High melting temps
  • Good electrical conductivity
  • Good thermal conductivity
  • Malleability
  • Ductility
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2
Q

Define metallic bonding

A

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

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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

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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

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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

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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

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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

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8
Q

What factors affect the strength of ionic bonding ?

A

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

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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

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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

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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

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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

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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.
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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

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15
Q

Define : covalent bond

A

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

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16
Q

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

A

A sigma bond

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17
Q

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

A

A sigma bond

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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

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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

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20
Q

What is bond length

A

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

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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

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22
Q

Define electronegativity

A

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

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23
Q

What is the general trend in electronegativity

A

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

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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
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25
Q

Define polar covalent bond

A

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

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26
Q

What is a discrete (simple) molecule

A

An electrically neutral group of two or more atoms held together by chemical bonds

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27
Q

How and when is a dative covalent bond formed?

A

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

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28
Q

How is a dative covalent bond often showed in displayed formula

A

By an arrow starting from the atom providing the electron pair and going towards the atom with the empty orbital

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29
Q

The electron pair repulsion theory states that…

A
  • 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
Q

When determining the shape of a molecule, how should you approach molecules with multiple bonds?

A

Treat each multiple bond as if it contained only one pair of electrons

31
Q

What is the shape and bond angle of a molecule with 2 bonded pairs ?

A

Linear
180 °

32
Q

What is the shape and bond angle of a molecule containing 3 bonded pairs

A

Trigonal planar
120 °

33
Q

What is the shape and bond angle of a molecule containing 4 bonded pairs ?

A

Tetrahedral
109.5 °

34
Q

What is the shape and bond angle of a molecule containing 6 bonded pairs ?

A

Octahedral
90 °

35
Q

What is the shape and bond angle of a molecule containing 2 bonded pairs and 1 lone pair ?

A

Non-linear
117.5 °

36
Q

What is the shape and bond angles of a molecule containing 3 bonded pairs and 1 lone pair

A

Trigonal pyramidal
107 °

37
Q

How do you calculate the bond angle for a molecule contains lone pair(s)?

A

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
Q

What is a dipole

A

The drift of bonded electrons towards the more electronegative element that results in a separation of charge

39
Q

What does the overall dipole (whether a molecule is polar/non-polar) depend on ?

A

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
Q

How is an instantaneous dipole created?

A

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
Q

How is the induced dipole created and how is that in turn responsible for the London forces?

A

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
Q

Describe & explain the trend in London forces as electrons per molecule increases

A

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
Q

Describe the trend in London forces as the shape and size of the molecule increases

A

The more points of interaction between molecules the greater the overall London force

44
Q

How do molecules with permanent dipoles interact and why is this interaction often weaker than that between instantaneous-induced dipoles?

A

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
Q

Describe hydrogen bonding through oxygen ?

A

All compounds containing an -OH group form intermolecular hydrogen bonds
The hydrogen bond forms between the oxygen atom of one molecule and the hydrogen atom of another (most important example is water)
There is also evidence for partial bond formation utilising the lone pair of electrons from the oxygen atom

46
Q

Describe hydrogen bonding through nitrogen

A

All compounds containing an N-H group can form intermolecular hydrogen bonds
Example : ammonia, NH3+

47
Q

Describe hydrogen bonding through fluorine

A

The only fluorine compound with intermolecular hydrogen bonding is hydrogen fluoride

48
Q

Define the hydrogen bond

A

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

49
Q

What are the reasons for the increase in boiling temp with increasing molecular mass in unbranched alkanes?

A

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

50
Q

How does branching affect the boiling points of alkanes

A

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

51
Q

What is the only significant intermolecular interaction in alkane molecules

A

London forces

52
Q

What are the significant intermolecular interactions between alcohol molecules ?

A

Contain an OH group so can form intermolecular hydrogen bonds in addition to London forces

53
Q

How does alcohols additional ability to form hydrogen bonds have an affect on the boiling temperature compared to alkanes

A

Hydrogen bonding provides an additional force of attraction that increases the energy required to separate the molecules

54
Q

What are the anamolous properties of water

A
  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 
55
Q

Explain the anomalous property of water : high melting and boiling temps

A

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

56
Q

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)

A

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

57
Q

Explain the anomalous property of water: the density of ice

A

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

58
Q

What are the conditions that must be met when choosing a suitable solvent ?

A

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

59
Q

When dissolving ionic solids, how is the energy required to seperate the ions in the solid supplied?

A

By the hydration of the ions

60
Q

Describe and explain the process of dissolving ionic solids in water

A

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’

61
Q

Which compounds can form hydrogen bonds with water

A

Alcohols
However the solubility of alcohols in water decreases with increasing hydrocarbon chain length as London forces predominate between alcohol molecules

62
Q

Which compounds cannot form hydrogen bonds with water

A

Non polar molecules such as alkanes
Halogenoalkanes (much more soluble in ethanol)
Some polar molecules

63
Q

Why do some polar molecules have limited solubility in water

A

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

64
Q

Describe a metallic lattice

A

Metallic lattices are composed of a regular arrangement of positive metal ions surrounded by delocalised electrons

65
Q

What are the four most common giant covalent lattices

A

Diamond
Graphite
Graphene
Silicon (IV) oxide

66
Q

Describe the bonding and properties of diamond

A

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

67
Q

Describe the bonding of graphite

A

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

68
Q

Describe the properties of graphite

A

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

69
Q

Describe the structure and bonding of graphene

A

One atom thick (pure carbon)
Carbon atoms are bonded in the same way as graphite

70
Q

Describe and explain the physical properties of molecular solids

A

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 (19 decks)

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