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Chemistry A- Level > Atomic structure. > Flashcards

Atomic structure. Flashcards

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1
Q

Define ionic bonding.

A

Definition: Ionic bonding is the electrostatic force of attraction between oppositely charged ions formed by electron transfer.

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2
Q

When is ionic bonding stronger. And how does this link to melting point.

A

Ionic bonding is stronger and the melting points higher when the ions are smaller and/ or have higher charges. E.g. MgO has a higher melting point than NaCl as the ions involved (Mg2+ & O2- are smaller and have higher charges than those in NaCl , Na+ & Cl- ).

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3
Q

Define a covalent bond.

A

A covalent bond is a shared pair of electrons.

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4
Q

Define a dative covalent bond and give three common examples drawing each one of them.

A

Definition: Shared pair of electrons where both electrons are donated by one atom. Common examples you should be able to draw that contain dative covalent bond are : NH4+, H3O+ (the hydroxonium ion), NH3BCl3.

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5
Q

How do dative covalent bonds act when thinking about the shapes of molecules. Hence give the shape of NH4+.

A

The dative covalent bond acts like an ordinary covalent bond when thinking about shape so in NH4+ the shape is tetrahedral.

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6
Q

What are the two ways of drawing a dative covalent bond. Draw in both of these ways NH3BF3

A

Draw the lewis structure (remembering square brackets and a charge if it’s an ion) or use arrows where the direction of the arrow goes from the atom that is providing the lone pair to the atom that is deficient.

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7
Q

Define metallic bonding.

A

Metallic bonding is the electrostatic force of attraction between the positive metal ions and the delocalised electrons.

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8
Q

Explain why Mg has a higher melting point than Na.

A

Mg has stronger metallic bonding than Na and hence a higher melting point. The Metallic bonding gets stronger because in Mg there are more electrons in the outer shell that are released to the sea of electrons. The Mg ion is also smaller and has one more proton. There is therefore a stronger electrostatic attraction between the positive metal ions and the delocalised electrons and higher energy is needed to break bonds.

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9
Q

The three main factors that affect the strength of metallic bonding are:

A

1) Charge on the Ion.
2. Number of delocalised electrons per atom (the outer shell electrons are delocalised)
The more delocalised electrons the stronger the bond
3. Size of ion.
The smaller the ion, the stronger the bond.

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10
Q

What two structures consist of covalent bonding.

A

1) Simple molecular crystal 2) Macromolecular crystal (giant covalent structures)

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11
Q

What type of structure consists of ionic bonding. Give examples.

A

Giant ionic lattice. NaCl , MgCl

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12
Q

Give examples of simple molecular structures (crystals
) what do they all have in common.

Give in layman’s terms what a simple molecule is.

A

Iodine Ice , Carbon dioxide, Water , Methane
They will all have some type of intermolecular bonding between molecules (doesn’t need to be the same type).

Only a few atoms held together by covalent bonds.

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13
Q

Give examples of macromolecular structures where the covalent bonds extend throughout the entire structure .

A

Diamond, Graphite ,Silicon dioxide, Silicon

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14
Q

What type of STRUCTURE does sodium have.

A

Giant metallic lattice.

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15
Q

Draw and name the shape of BeCl2.

A

Linear - 180 degree bond Angle. Single bonds.

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16
Q

Draw and name the shape of HNC.

A

Linear . Nitrogen forms its typical triple bond with carbon and a single bond with hydrogen. Remember the number of shared electrons in a bond doesn’t affect shape a bond it a bond.

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17
Q

Draw and name the shape of SO3

A

Trigonal Planar. Sulphur forms three double bonds with oxygen and so ends up with am extended octet of 12 electrons.

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18
Q

Draw and name the shape of NO3-

A

Trigonal planar. The Nitrogen forms 2 single bond and one double bond with the oxygen. Need square brackets with -1 charge on the outside. Total of 8 electron in the extended octet.

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19
Q

Draw and name the shape of CO32-

A

Trigonal planar. The carbon will form one double bond and also 2 single bonds. Remember the square brackets and the -1 charge on the outside.

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20
Q

Draw and name the shape of ClO4- and state the special property of chlorine in this example.

A

The shape is tetrahedral where there is 3 double bonds and one single bond ( hence the -1 charge from not getting rid of the electron ). The property of the chlorine is that it has a ridiculously extended octet of 14 electrons.

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21
Q

Draw and name the shape of SO42-.

A

The shape is just tetrahedral and can be drawn just like the sulphuric acid molecule but without the H’s on the single bond oxygens.

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22
Q

Draw and name the shape of NH4+.

A

Tetrahedral with 4 double bonds.

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23
Q

What element will typically form trigonal pyramidal structures and name two possible examples. What form will they almost defo come in as a give away.
Hint this is this same form as trigonal planar.

A

Generally will be a group 5 element bonded to something that only needs one electron so that when they bond you get 2 electrons ( 1 pair) that is left over.
Examples include: NCl3 or AsCl3 or PF3
all of the above examples are group 5 elements bonding with halogens which is typical of creating a trigonal pyramidal shape. Always in XY3 form or X3Y
There is some exceptions with Oxygen that also work.

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24
Q

Draw and name the structure of OF2.

A

Bent 2 single bonds.

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25
Q

Draw and name the structure of H2S .

A

Bent , 2 single bonds.

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26
Q

Draw and name the structure of SCl2.

A

bent , 2 single bonds.

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27
Q

Draw and name the structure of PCl5

A

Trigonal Bipyramidal - in 3D RHS should have 3 straight lines and the LHS should have a wedge and a dotted line.

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28
Q

Draw and name the structure of SF6.

A

Octehedral - For 3D straight lines vertically up an vertically down and then two diagonal up wedges and two diagonal down dotted lines.

29
Q

Draw and name the structure of I3 - (iodine)

What do you need to remember about the numbers , and what is a helpful way to think about the lone pairs.

A

It is linear in shape. It has an extended octet with two singly bonded Iodines and 3 lone pairs since the negative charge adds one on. The lone pairs arrange themselves so they even out (sort of in there own trigonal planar structure) with 2 on one side and 1 own the other and so you get this linear shape. SEE PICTURE .
Only 2 iodines bond to the central one not three.

30
Q

Draw and name the structure of KrF2 and hence what shape will a noble gas and a halogen be when ions ions this XY2 form.

A

It is linear since as with I3- if you have 3 lone pairs and two bonding pairs you lone pairs will arrange themselves in a trigonal planar way that evens out so the bonding pairs are arranged 180 degrees apart in a linear fashion.

31
Q

Draw the shape of ClF3 give a brief explanation of this shape and state also the special property of the chlorine atom here. What standard shape of molecule is this a variation of and hence what can we conclude is the shape of two halogens bonded together.

A

This is a simple variation of a 5 pair total trigonal bi-pyrimidal structure. There is 2 lone pairs and three bonding pairs (unlike the inverse which gives way to a linear structure). You end up with a T-shaped molecule. Both lone pairs with their little hats come out of one side and the three bonding pairs come out of the other side in a t-shape.

32
Q

Give the general description for each type type of bonding and explain why what you have said for metals is true (in terms of malleability).

A

Ionic crystals are crystalline solids
Simple molecular compounds aremostly gases and liquids.
Macromolecular compounds are always solids.
Compounds that have a giant metallic lattice structure are shiny metals and malleable sincethe positive ions in the lattice are all identical. So the planes of ions can slide easily over one another
-attractive forces in the lattice are the same whichever ions are adjacent.

33
Q

Give the solubility of the 4 different types of crystals.

A

Ionic compounds have good solubility
Simple molecular compounds will have poor solubility.
Both macromolecular and metallic compounds will be insoluble.

34
Q

Draw and name the structure of CLF4-.

A

g

35
Q

Draw and name the structure of H3O.

A

g

36
Q

What can a mass spectrometer be used for.

A

Determining the all the isotopes present in a sample and hence identify elements or find the Mr of an organic substance from the molecular ion peak.

37
Q

State why mass spectrometry needs to be done under a vacuum.

A

So the ions dont collide with air particles and also so that air particles don’t ionise and be picked up by the detector.

38
Q

Describe fully electrospray ionisation. Give an equation where X is ionised.

A

A sample is dissolved in a polar volatile solvent.
Injected through a fine hypodermic needle.
Tip of the needle is connected to the positive end of a power supply with a High voltage.
At the needles tip the sample molecule M , will gain a proton forming MH+.
The solvent evaporates.

M gas + H+ INTO MH+ gas.

39
Q

Describe fully electron impact ionisation.

A

A vaporised sample is bombarded with high energy electrons from an electron gun. This nocks out an outer electron forming a charged Ion.
X gas INTO X+ gas add an electron.

40
Q

Describe when each method of ionisation should be used.

A

You generally see larger organic molecules being done with electropsray since it is gentler and less likely to cause fragmentation. Electron impact is used for elements with lower Mrs for example single elements that need to have isotopes determined as opposed to finding the Mr of am organic molecule that would fragment with electron impact.

41
Q

Describe fully what happens during the acceleration stage.

A

Positive ions are accelerated by an electric field to constant velocity.

42
Q

Describe what happens in the flight tube and why this means mathematically ions can be separated.

A

Because all the particles have the same kinetic energy, the velocity of each particle is dependent o its mass. Lighter particles will have a greater velocity and vice versa.

43
Q

Give the units that should be used for any TOF calculations.

A

Ke is in J Not KJ
Mass should be In Kg not grams this is most important as you need to divide your Avogadro division by a thousand before using it.
Velocity is In meters per second.
Distance is in meters.

44
Q

What two variables does the TOF spectrometer measure.

A

Percentage abundance and MZ

45
Q

What is the MZ value of Mg2+

A

12

46
Q

What should you give when asked to give the species responsible for a peak.

A

Give the species with the mass number or MZ on the LHS and the CHSRGE included on the right eg 24Mg+
small 24 in top left.

47
Q

What is a practical application of TOF mass spectrometry.

A

Included in space probes where elements on other planets have different compositions of isotopes.

48
Q

For 6 marks compare how you could measure the Mr of butane using electro spray and electron impact methods of ionisation.

A

Write quick.

49
Q

Explain the progression of the atomic model.Give used for each.

A

Daltons indivisible atoms.

Rutherford.

Bohr .

Shrodinger.

50
Q

Draw the shapes of the s p and d orbital.

A

S is a sphere
P is a figure of 8 loop
d is a double figure of 8 loop.

51
Q

What does an orbital represent.

A

The region of space with a 95 percent probability of finding an electron around the nucleus. Each orbital has its own unique shape.

52
Q

Draw a spin diagram for fluorine and state what the box and the arrow represent. Comment on the vital method for writing out spin diagrams.

A

1 s one box
2s one box
2p three boxes. going upward to show increasing energy from bottom to top.
The box represents an orbital which can only ever hold two electrons of opposite spin.

53
Q

Give the electronic configuration of Cr3+ and Cu2+ , state what needs to be done when writing electron configurations of d block elements and what is special about the two examples above.

A

Write out and check from sheet.

54
Q

Define first ionisation energy and state the three factors that will affect it.

A

Enthalpy change when on mole of gaseous atoms form one mole of gaseous atoms to form one moles of gaseous ions with a single positive charge.

55
Q

Why are successive ionisation energies always larger.

A

Removing an electron from a more positive species.

56
Q

Give a word perfect answer how you could conclude that an element was In group 2 based on successive ionisation enthalpies.

A

There is a big jump between the 2nd and third ionisation energies which mean the third electron must have been removed from an electron shell closer to the nucleus with less shielding so has a larger ionisation enthalpy.

57
Q

Why does helium have the greatest ionisation energy.

A

`First electron in shell closest to nucleus.Has one more proton than hydrogen.

58
Q

Why do first ionisation energies decrease down the group.

A

Outer electrons in shells further away and better shielded. Weaker attraction.

59
Q

Why is there there a general increase in first ionisation energy across a period.

A

More protons for the same shielding effect as electron are added to the same shell the effective charge is greater.

60
Q

Why does Na have a much lower ionisation energy than neon.

A

Simply since Na will have its outer electron in the 3s orbital so is further from the nucleus and better shielded so easier to remove.

61
Q

Explain the small drops ionisation enthalpies across period 3 using spin diagrams.

A

Look at sheet.

62
Q

Draw a graph of second ionisation enthalpies across period 3 and explain which element now has a larger second ionisation enthalpy , helium or lithium.

A

The group one elements are now at the peak of the graph as there second electron is being removed from a shell closer to the nucleus. Sodium has the highest second ionisation energy of them all. Li has more protons so since its electron is now being removed from the same closest shell to the nucleus as helium it is a question of what has more protons and lithium does.

63
Q

Define relative atomic mass.

A

Average weighted mass of an atom of an element taking into account its naturally occurring isotopes relative to one twelfth the mass of carbon 12.

64
Q

What should you quote when asked to give the species responsible for a given peak.

A

Always quote the ion or molecular ion. 79Br + for example for electron impact. Or XH+ for elctrospray.
Maybe C3H10 + if electron impact used with fragmentation.

65
Q

Draw the mass spectra for Cl2 and for Br2 and label the species responsible for all the peaks. Draw the peaks caused by the diatomic molocules.

A

For chlorine:
Peak at 7O for Cl35Cl35+
Peak at 72 For Cl35Cl37+
and one at 74 for Cl37Cl37+ staircase down

Bromine has one big middle one at 160 due to then double affect then smaller ones at 158 and 162.

66
Q

What must you think about when asked to determine the Mr of a molecule when electro spray ionsisation has been used. Say you had one big peak at 521.1 what is the Mr of the species if electro spray ionisation was used.

A

This peak is the Mr plus one for the H+ so this example would be 520.1

67
Q

Explain what causes peaks in different types of mass spectrometer and what the general look is for different types of spectrum.

A

s

68
Q

Explain the detection stage in TOF.

A

Ions generate current when they reach the detector by flow off electrons from the detector to the positive ions.
Size of the current is proportional to abundance.

69
Q

Draw spin diagrams for both Magnesium and aluminium and show what complete removal during ionisation would look like.

A

s

Chemistry A- Level (18 decks)

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