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Chemistry Year 12/13 > Module 2.2 - Electrons, Bonding and Structure > Flashcards

Module 2.2 - Electrons, Bonding and Structure Flashcards

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1
Q

What does the principal quantum number, n, show?

A

The shell that the electrons occupy (larger n = further from nucleus)

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2
Q

How many electrons can fit in the 1st shell?

A

2

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3
Q

How many electrons can fit in the 2nd shell?

A

8

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4
Q

How many electrons can fit in the 3rd shell?

A

18

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5
Q

How many electrons can fit in the 4th shell?

A

32

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6
Q

How many electrons can an s orbital hold?

A

2

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7
Q

How many electrons can an f orbital hold?

A

2!

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8
Q

What is the shape of an s orbital?

A

spherical

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9
Q

What is the shape of a p orbital?

A

dumbbell

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10
Q

How many electrons can a d sub-shell hold?

A

10

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11
Q

How many electrons can a p sub-shell hold?

A

6

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12
Q

How many electrons can an f sub-shell hold?

A

14

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13
Q

Why does the 4s orbital fill before the 3d orbital?

A

Is a slightly lower energy level

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14
Q

How do atomic orbitals fill?

A
  • from lowest energy level upwards
  • each energy level must be full before next higher energy starts to fill
  • 2p orbitals are filled singly before pairing starts at oxygen
  • paired electrons have opposite spins
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15
Q

Between what type of elements does ionic bonding usually occur?

A

metal and a non metal

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16
Q

What is meant by the term ‘ionic bond’?

A

Electrostatic attraction between oppositely charged ions

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17
Q

What is meant by the term ‘covalent bond’?

A

The electrostatic attraction between a shared pair of electrons and the nuclei of the bonded atoms

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18
Q

Between what types of elements does covalent bonding usually occur?

A

2 non metals

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19
Q

What is meant by the term ‘metallic bond’?

A

Strong electrostatic attraction between cations (positive ions) and delocalised electrons

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20
Q

How does an ionic bond form?

A

Elections transferred from metal to non metal to form oppositely charged ions, bonded by electrostatic attractions. Each ion is surrounded by the oppositely charged ions, and ions attract each other from all directions forming a giant ionic lattice

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21
Q

What are the properties of ionic compounds?

A
  • high melting and boiling points
  • electrical conductivity (when aqueous or molten)
  • solubility
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22
Q

Why do ionic substances have a high melting and boiling point?

A
  • solid at room temp
  • lots of energy needed to break strong electrostatic bonds between oppositely charged ions in solid lattice
  • the greater the charge, the stronger the electrostatic forces between the ions so more energy to break
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23
Q

Compare the electrical conductivity of ionic substances when solid and when they’re aqueous/molten.

A
  • solid: ions held in fixed position and can’t move so doesn’t conduct electricity
  • aqueous/molten: ions are free to move so can now conduct electricity
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24
Q

Describe the solubility of ionic substances.

A
  • soluble in polar solvent e.g. water as contain substances w polar bonds
  • water surrounds each ion to form a solution; slight charges within polar substance are able to attract charged ions in giant ionic lattice so lattice disrupted and ions pulled out of it
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25
Q

What is average bond enthalpy?

A

Measure of covalent bond strength, the higher the number, the stronger the covalent bond so more energy needed to break them

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26
Q

What does the direction of the arrow show when representing a dative covalent bond?

A

Direction which the electron pair has been donated

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27
Q

What is a dative bond also known as?

A

coordinate bond

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28
Q

Give 2 examples of where a dative bond is seen.

A
  • ammonium ion

- oxonium ion, H3O+ (2 covalent and 1 dative covalent and one lone pair)

29
Q

What is the octet rule?

A

That atoms bond to get 8 electrons in outermost shell

30
Q

Why is the octet rule not always followed?

A
  • not enough electrons to reach octet

- more than 4 electrons may pair up in bonding (expansion of the octet)

31
Q

Give one example of when there aren’t enough electrons from the octet.

A

BF3

  • 3 covalent bonds can be formed
  • each of boron’s 3 outer electrons paired
  • F achieves octet but boron doesn’t have enough electrons
32
Q

Which group’s elements can have expansion of the octet?

A
  • 15: 3 or 5 covalent bonds depending on how many electrons used in bonding
  • 16: 2, 4 or 6 depending on how many used in bonding
  • 17: 1, 3, 5 or 7 depending on how many used in bonding
33
Q

Give an example of a molecule where expansion of the octet occurs.

A

SF6

  • 6 outer shell electrons
  • 6 covalent bonds form so now 12 outer shell electrons (no longer obeys Octet Rule)
  • fluorine gets the octet
34
Q

How could the octet rule be modified?

A
  • unpaired electrons pair up

- max no. of electrons that can pair is equivalent to no. of electrons in outer shell

35
Q

What are the 2 types of covalent structure?

A
  • simple molecular lattice

- giant covalent lattice

36
Q

What is the structure of simple molecular structures?

A
  • atoms within each molecule help together by strong covalent bonds
  • weak intermolecular forces holding different molecules together
37
Q

What are the properties of simple molecular structures?

A
  • low melting and boiling points (weak intermolecular forces so little energy needed to break)
  • non conductors of electricity as no charged particles free to move
  • soluble in non polar solvents like hexane (weak London forces able to form between covalent molecules and these solvents so molecular lattice breaks down and substance dissolves)
38
Q

What are the properties of giant covalent structures?

A
  • high melting and boiling points (high temps needed to break strong covalent bonds within lattice)
  • non conductors of electricity as no free charged particles
  • insoluble in polar and non polar solvents (covalent bonds in lattice are too strong to be broken by either type of solvent)
39
Q

What is the electron repulsion theory?

A
  • all electrons have a negative charge so repel one another

- shape adopted will be the shape that allows electrons to be as far apart as possible

40
Q

What shape forms when theres one electron pairs bonding around a central atom, like in hydrogen gas?

A

linear

41
Q

What shape forms when there are 2 bonding regions and no lone pairs around the central bonding atom, like in carbon dioxide?

A
  • linear

- 180º bond angle

42
Q

What shape forms when there are y bonding regions and no lone pairs around the central bonding atom, like in boron trifluoride?

A
  • trigonal planar

- 120º bond angle

43
Q

What shape forms when there are 4 bonding regions and no lone pairs around the central bonding atom, like in methane?

A
  • tetrahedral

- 109.5º bond angle

44
Q

What shape forms when there are 5 bonding regions and no lone pairs around the central bonding atom, like in phosphorus pentachloride?

A
  • trigonal bipyramidal

- 90º and 120º bond angles

45
Q

What shape forms when there are 6 bonding regions and no lone pairs around the central bonding atom, like in sulphur hexafluoride?

A
  • octahedral

- 90º bond angle

46
Q

What shape forms when there are 2 bonding regions and 2 lone pairs around the central bonding atom, like in water?

A
  • non linear

- 104.5º bond angle

47
Q

What repels more, lone pairs or bonded pairs of electrons?

A
  • lone pairs, reduce bond angle by 2.5º

- lone pair/lone pair>bonded pair/lone pair>bonded pair/bonded pair

48
Q

What shape forms when there are 3 bonding regions and 1 lone pair around the central bonding atom, like in ammonia?

A
  • pyramidal

- 107º bond angle

49
Q

What is meant by the term ‘electronegativity’?

A

The ability of an atom to attract the bonding electrons in a covalent bond

50
Q

Where does electronegativity increase towards on the Periodic Table?

A

top right - fluorine

51
Q

What does the Pauling scale measure?

A

the electronegativity of an atom

52
Q

How does a dipole form?

A
  • covalent bond where 1 atom is more electronegative than the other
  • more electronegative atom has more negative dipole
  • forms a polar covalent bond
53
Q

How can a molecule with polar bonds, such as tetrachloromethane, be non polar?

A
  • C-Cl bond is polar
  • molecule is symmetrical
  • dipoles act in different directions and cancel each other out
54
Q

Is water polar and why?

A
  • yes
  • has an overall dipole
  • not symmetrical so dipoles don’t cancel
  • is non linear
55
Q

Which are the strongest: hydrogen bonds, permanent dipole-permanent dipole interactions or London forces?

A
  • H bonds
  • then permanent dipole-permanent dipole interactions
  • then London forces
56
Q

How does a permanent dipole-induced dipole interaction form?

A
  • permanent dipole on one molecule
  • when near a non polar molecule it can cause electrons in shells to shift slightly as -ve repels -ve or attracts +ve
  • molecule with permanent dipole has induced a dipole on the other molecule
57
Q

How does a permanent dipole-permanent dipole interaction form?

A
  • molecules w permanent dipoles attracted to other molecules w permanent dipoles
  • opposite delta charges attracted to one another e.g. in HCl
58
Q

What 2 types of interactions come under the bracket of permanent dipole-dipole interactions?

A
  • permanent dipole-induced dipole interactions

- permanent dipole-permanent dipole interactions

59
Q

How are London dispersion forces formed?

A
  • caused by constant random movement of electrons in atoms’ shells. Movement unbalances distribution of charge within electron shells (electron density moves from side to side)
  • at any moment there’ll be an instantaneous dipole across molecule
  • instantaneous dipole induces a dipole in neighbouring molecule, which induces further dipoles on their neighbouring molecules
  • small induced dipoles attract one another, causing weak intermolecular forces (London dispersion forces) or instantaneous dipole-induced dipole forces
60
Q

How does number of electrons affect the strength of London forces?

A
  • size of force increases w more electrons

- larger induced dipoles so greater attraction between molecules

61
Q

Why do molecules with van der Waals’ forces have a low boiling and melting point?

A

Weak intermolecular bonds that need to be broken

62
Q

Describe the trend in boiling point of the noble gases as you go down the group.

A

Increases as more electrons so stronger London forces

63
Q

What elements can a hydrogen bond form between?

A

O-H
N-H
F-H

64
Q

Give examples of substances with hydrogen bonding.

A
  • ammonia

- water

65
Q

What does a hydrogen bond form between?

A

hydrogen atom and lone pair of electrons of O, N or F

66
Q

What properties of water result from its hydrogen bonding?

A
  • ice is less dense than water
  • high melting and boiling point for its size
  • high surface tension and viscosity (why insects can walk on water)
67
Q

Explain why ice can float on water.

A
  • when ice forms, H2O molecules arrange to orderly pattern and H bonds form between molecules (occurs in liquid phase too but not as often as molecules move past each other hence overcome these bonds)
  • ice has open lattice w H bonds holding H2O molecules apart
  • when ice melts, rigid H bonds collapse, allowing H2O molecules to move closer
  • so ice less dense than water hence why ice floats
68
Q

Why does water have a higher melting and boiling point than expected?

A
  • H bonds much stronger than other intermolecular forces
  • extra strength of these forces has to be overcome to melt/boil H2O. Results in H2O having higher melting/boiling points than expected if H bonds weren’t present
  • other group 16 hydrides e.g. SiH2 show same structure of water but have no H bonding, so much lower melting and boiling point
69
Q

What is meant by the term ‘hydrogen bond’?

A

A strong permanent dipole-permanent dipole attraction between an electron efficient hydrogen atom on one molecule, and a lone pair of electrons on a highly electronegative (N, O, F) atom on a different molecule

Chemistry Year 12/13 (18 decks)

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