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Organic Chem AS Mock > TOPIC 1: Electron, Bonding and Structure > Flashcards

TOPIC 1: Electron, Bonding and Structure Flashcards

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1
Q

Explain the different boiling points of NH3, F2 and BR2. (5)

A
  • NH3 has hydrogen bonding b/w molecules
  • F2 & Br2 have induced d.p-d.p forces
  • intermolecular forces in Br2 stronger than in F2

^ bcs Br2 has more electrons

  • Hydrogen bonding is stronger than the induced d.p-d.p forces in F2
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2
Q

State and explain two anomalous properties of ice caused by hydrogen bonding. (4)

A
  • ice is less dense than water
    ^ ice has an open lattice
  • ice has a relatively high melting point
    ^ hydrogen bonds are relatively strong
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3
Q

Describe and explain the electrical conductivity of sodium oxide and sodium in their solid and molten states. (5)

A
  • sodium conducts in solid and molten state
    ^ sodium has delocalised electrons in both states
  • sodium oxide conducts in molten state but not solid
    ^ molten sodium oxide has mobile ions
    ^ solid sodium oxide has immobile ions
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4
Q

Why is Na3PO4 described as a salt of H3PO4. (1)

A
  • the hydrogen ions are replaced by sodium ions
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5
Q

Suggest why PH3 has a lower boiling point than NH3. (1)

A
  • the intermolecular forces in PH3 are weaker
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6
Q

Define: a dative covalent bond

A
  • both electrons have been donated by one atom
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7
Q

Explain why the student might expect the H-N-H bond angle to be larger in H3NBF3 than in NH3. (3)

A
  • NH3 has 3 bonding pairs & 1 lone pair
    of electrons
  • H3NBF3 has 4 bonding pairs of
    electrons
  • lone pair on Nitrogen now becomes a
    bonding pair
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8
Q

Explain why there is a difference in the melting points of K, KBr and H2O. (6)

A
  • In K, there’s electrostatic attraction
    b/w cations & electrons
  • In KBr, there’s electrostatic attraction
    b/w oppositely charged ions
  • KBr has ionic bonding
  • In H2O, there’s hydrogen bonding b/w
    molecules
  • strength of forces: (bonding)
    ionic > metallic > hydrogen
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9
Q

Define: electronegativity

A
  • the ability of an atom to attract electrons
  • in a covalent bond
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10
Q

Molecules of BF3 contain polar bonds, but the molecules are non polar.
Explain why. (2)

A
  • BF3 is symmetrical
  • the dipoles cancel out
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11
Q

Explain why a CH2Cl2 molecule is polar. (1)

A
  • the dipoles do not cancel out
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12
Q

SbCl3 molecules are polar.
Explain. (2)

A
  • there’s a difference in
    electronegativity
  • molecules are not symmetrical and
    dipoles do not cancel out
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13
Q

Describe how Van der Waals’ forces arise. (3)

A
  • uneven distribution of electrons
  • creates a temporary dipole
  • causing induced dipoles in
    neighbouring molecules
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14
Q

Suggest why there are no other intermolecular forces in solid sulfur. (1)

A
  • no polar bonds
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15
Q

Explain why a molecule of SF6 has an octahedral shape. (2)

A
  • sulfur has six bonded pairs
  • electron pairs repel
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16
Q

Explain why a solution of copper(II) nitrate conducts electricity. (1)

A
  • ions are mobile
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17
Q

State and explain the trend in the boiling points of chlorine, bromine and iodine. (3)

A
  • b.p. increases down the group
  • stronger intermolecular forces
  • more energy needed to break these
    intermolecular forces
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18
Q

Why would astatine be expected to react similarly to other halogens. (1)

A
  • same number of outermost electrons
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19
Q

Explain why a solution of Copper(II) nitrate conducts electricity. (1)

A
  • ions are mobile
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20
Q

Explain how ionic bonding holds the particles together in an ionic compound. (1)

A
  • strong electrostatic attraction between positively charged ions & negatively charged ions
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21
Q

Describe the bonding process that forms lithium halides, with respect to electrons. (3)

A
  • Li donates one electron
  • The halogen accepts one electron
  • Li forms a 1+ ion
    & Cl forms a 1- ion
22
Q

Describe the type of bond that forms b/w the boron triflouride and ammonia molecules. (2)

A
  • a dative covalent bond
  • the N provides both electrons for the bond
23
Q

Describe the bonding in an (S2)2- ion. (1)

A
  • 2 sulfurs attached by a covalent bond
    ^ each with a negative charge
24
Q

Explain why silver chloride has such a high melting point. (3)

A
  • giant ionic structure
  • strong electrostatic attraction b/w oppositely charged ions

^ which requires large amounts of energy to overcome

25
Q

Explain why Al2Cl6 doesn’t conduct when molten but Al2O3 does. (4)

A

Al2Cl6:
- covalent bonding
- therefore no free charges when molten, so not an electrical conductor

Al2O3:
- has an ionic lattic
- therefore there’s mobile ions in the molten state

26
Q

Describe the bonding in calcium carbide, CaC2. (4)

A
  • (C2)2- is a diatomic ion
  • triple bond b/w carbons
  • Ca2+ ion
  • ionic bond b/w Ca2+ and (C2)2-
27
Q

Explain why ionic compounds have high melting and boiling points (2)

A
  • Strong electrostatic attractions between oppositely charged ions.
  • High temperature needed to provide sufficient energy to overcome the attractions.
28
Q

Explain why ionic compounds dissolve in water (2)

A
  • Polar water molecules are attracted towards ions on the surface of the ionic lattice. Water molecules bond to the ions, weakening and breaking them.
  • Ions become surrounded by water molecules and break free from the lattice
29
Q

Define: covalent bonding

A

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

30
Q

Define: orbital

A

A region around the nucleus that can hold up to two electrons with opposite spins

31
Q

Define: isoelectronic

A

The same number of electrons

32
Q

State the maximum number of electrons that can be held in each of the first three shells of an atom. (3)

A

1st = 2
2nd = 8
3rd = 18

33
Q

State how many orbitals there are in a p-subshell and how the electrons are arranged if the subshell is full (3)

A
  • 3 atomic orbitals
  • each orbital has 2 electrons
  • with opposite spins
34
Q

Explain why both G1 & G2 are known as s-block elements (1)

A

The outer electrons are s-subshell electrons for all

35
Q

Explain the strength of the ionic bond in sodium chloride (2)

A
  • strong electrostatic attraction
  • b/w Na+ and Cl- ions
36
Q

Explain why it’s easier to use aqueous MgF2 in a lab setting than the molten version (2)

A
  • High melting point
  • as strong electrostatic attraction b/w oppositely charged ions in all directions
37
Q

Compare the binding in phosphorus trichloride and ammonia (4)

A
  • both undergo covalent bonding
  • phosphorus atom and nitrogen atom are both central atoms
  • each have 1 lone pair
  • 3 shared pairs of electrons
38
Q

Describe how the shape of the periodic table is linked to the electronic structure (6)

A
  • elements in same period have same number of filled shells
  • elements in same group have same number of valence electrons
  • s-block includes G1 & G2 as it takes only 2 electrons to completely fill s-subshell
  • p-block includes G3-0 as it takes 6 electrons to completely fill the p-subshell
  • d-Block includes transition metals as it takes 10 electrons to completely fill d-subshell
  • d-block begins on period 4 as d-orbitals have higher energy than 4s orbitals
39
Q

Suggest why there are three possible p-subshells but only one possible s-subshell in an atom (3)

A
  • s-orbitals are spherical so multiple subshells not possible
  • p-orbitals are propeller shaped so 3 p-orbitals would not overlap significantly
40
Q

Justify why hydrogen is positioned in the middle of the periodic table and not apart of G1 (3)

A
  • very diff physical properties to G1 metals
  • and chemical properties
  • despite outer shell containing 1 s-subshell electron only
41
Q

Explain why giant ionic structures have high melting points (2)

A
  • strong electrostatic attraction between oppositely charged ions
  • large amounts of energy required to overcome them
42
Q

Explain why sodium bromide has a higher melting point than sodium and sodium iodide (6)

A

Stage 1:
- Na has metallic bonding & a giant structure
- there’s attraction b/w positive nucleus and delocalised electrons in Na

Stage 2:
- ionic bonding in NaBr & giant structure
- there’s attraction b/w + and - ions in NaBr

Stage 3:
- ionic bonds are stronger than metallic bonds
- stronger attraction b/w the + and - ions in NaBr than NaI
- since Br- ion is smaller than I- ion

43
Q

Why is an arrow used to represent a
N ➡️ H bond? (1)

A

To show both electrons come from nitrogen

44
Q

Suggest how methanol and methanethiol could be separated (1)

A

(Fractional) distillation

45
Q

Define: ionic lattice (2)

A
  • Repeating pattern
  • of oppositely charged ions
46
Q

State whether the following conduct electricity when solid or molten: (5)

aluminium
aluminium fluoride
boron tribromide

A

Aluminium:
- conducts in solid and molten stated
- has delocalised electrons

Aluminium flouride:
- conducts when molten because it has mobile ions
- doesn’t conduct when solid (ions fixed in position in an ionic lattice)

Boron tribromide:
- does not conduct in solid and molten states (no mobile ions)

47
Q

Explain how the structure and bonding in bromine account for its relatively low melting point (3)

A
  • forces b/w molecules
  • which are induced dipole-dipole forces
  • are weak so overcome easily by increased kinetic energy
48
Q

Predict the type of structure and bonding of SO2 and MgO & explain the difference in their melting points (4)

A
  • MgO = giant ionic
  • SO2 = simple molecule
  • ionic bonds in MgO much stronger than intermolecular bonds in SO2
  • ionic bonds in MgO require more energy to overcome
49
Q

Describe the relative energies of the 2s orbitals and each of the three 2p orbitals in a nitrogen atom (2)

A
  • p-orbitals have greater energy than
    s-orbitals
  • three p-orbitals have equal energy
50
Q

What are the 3 rules for filling electron orbitals? (3)

A
  • Hund’s rule (each orbital must be singlely occupied before being paired up)
  • Each orbital can hold up to 2 electrons with opposite spins
  • Lowest energy orbitals must be filled first
51
Q

Why is hyrogen H+

A

fills 1s orbital

Organic Chem AS Mock (3 decks)

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