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BONDING Flashcards

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1
Q

State the structure of, and bonding in, silicon dioxide. Other than a high melting point, give two physical properties of silicon dioxide that are characteristic of its structure and bonding.
(4)

A

Structure: Macromolecular / giant molecule / giant covalent
Bonding: Covalent / giant covalent

Physical Properties:
Any two from: Hard/ Brittle / not malleable Insoluble Non conductor

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2
Q

Give the formula of the species in a sample of solid phosphorus(V) oxide. State the structure of, and describe fully the bonding in, this oxide.
(4)

A

Formula: P4O10
Structure: Molecular
Bonding: Covalent / shared electron pair
van der Waals’ / dipole–dipole forces between molecules

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3
Q

Explain why magnesium has a higher melting point than sodium.

(3)

A

more protons (1) (or Mg2+ more charge than Na+) attracts delocalised (or bonding) electrons more strongly (1) Metallic bonding is stronger

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4
Q

Why do diamond and graphite both have high melting points?

(2)

A

macromolecular (1) strong covalent bonds (1)

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5
Q

Why is graphite soft?
(2)

A

Planes (1) weak (bonds) forces between planes (1)

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6
Q

The table below shows the melting points of some sodium halides.

                           NaCl  NaBr  NaI

Melting point /K 1074 1020 920

Suggest why the melting point of sodium iodide is lower than the melting point of sodium bromide.

(1)

A

Iodide / I– bigger (ion) (so less attraction to the Na+ ion)

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7
Q

Fluorine reacts with bromine to form liquid bromine trifluoride (BrF3). State the type of bond between Br and F in BrF3 and state how this bond is formed.
(2)

A

Covalent
Shared pair(s) of electrons / one electron from Br and one electron from F

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8
Q

When H+ ions react with H2O molecules, H3O+ ions are formed.

Name the type of bond formed when H+ ions react with H2O molecules. Explain how this type of bond is formed in the H3O+ ion. (2)

A

Dative (covalent)/coordinate

(Lone) pair/both electrons/two electrons on O(H2) donated (to H+) OR pair/both electrons come from O(H2)

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9
Q

Explain why the AsCl4+ ion has a bond angle of 109.5* (As in group 5)

A

There are 4 bonds or 4 pairs of electrons (around As)
(Electron pairs / bonds) repel equally

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10
Q

what is the meaning of the symbol δ+

A

Electron deficient

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11
Q

Methanol has the structure

Explain why the O–H bond in a methanol molecule is polar.

A

Oxygen more electronegative than hydrogen

causes higher e– density round oxygen atom / causes Hδ+ Oδ–

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12
Q

Suggest why the electronegativity of the elements increases from lithium to fluorine.(2)

A

More protons / bigger nuclear charge

Same or similar shielding

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13
Q

Name the crystal type which describes an iodine crystal

A

molecular

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14
Q

Describe the bonding in a crystal of iodine.

A

Van der Waals’ between molecules (1)

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15
Q

In terms of the intermolecular forces involved, suggest why

(i) hydrogen fluoride requires more heat energy for melting than does hydrogen chloride,

A

H-Bonding in HF (1)
(dipole-) dipole in HCl (1)
H-bonding is stronger than dipole-dipole or V.dW (1)

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16
Q

In terms of the intermolecular forces involved, suggest why
hydrogen iodide requires more heat energy for melting than does hydrogen chloride

A

HI bigger molecule than HCl (1)
OR more electrons
Therefore the forces between HI molecules are stronger (1)

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17
Q

Explain how the strongest type of intermolecular force in liquid HF arises.

A

large difference in electronegativity between H and F
δ+H-Fδ- dipole created or dipole clearly implied
attraction/bond formed between δ+H and lone pair on F

18
Q

Explain the trend in the boiling points of the hydrogen halides from HCl to HI.
HCl HBr HI

Boiling point / K 188 206 238

A

van der Waals’ / induced/temporary dipole-dipole attractions

increase with the increasing Mr / size / mass / number of electrons (in the hydrogen halides)

19
Q

Silicon tetrafluoride (SiF4) is a tetrahedral molecule.
Deduce the type of intermolecular forces in SiF4
Explain how this type of intermolecular force arises and why no other type of intermolecular force exists in a sample of SiF4

A

Van der Waals forces
(Uneven distribution of electrons in) one molecule induces dipole in neighbouring/another molecule
symmetrical molecule / dipoles cancel

OR
no hydrogens bonded to F (N or O), therefore no hydrogen bonding

20
Q

The boiling points of fluorine and hydrogen fluoride are –188 °C and 19.5 °C respectively. Explain, in terms of bonding, why the boiling point of fluorine is very low.

A

VdW / van der Waals forces between molecules

IMF are weak / need little energy to break IMF / easy to overcome IMF

21
Q

The bond angle around the oxygen atom in methanol is slightly smaller than the regular tetrahedral angle of 109.5°

Explain why this bond angle is smaller than 109.5°

A

Idea that lone pairs have greater repulsion than bonding pairs

22
Q

Explain how permanent dipole-dipole forces arise between hydrogen chloride molecules.

A

Difference in electronegativity leads to bond polarity

(dipoles don’t cancel therefore the molecule has an overall permanent dipole) and there is an attraction between ∂+ on one molecule and ∂− on another

23
Q

Van der Waals’ forces exist between all molecules.

Explain how these forces arise.

A

Electron movement in molecule causes temporary dipole/uneven distribution of electrons

Induces a dipole in another molecule

δ+ attracts δ- in different/adjacent molecules

24
Q

Phosphine (PH3) has a structure similar to ammonia.

In terms of intermolecular forces, suggest the main reason why phosphine is almost insoluble in water.

A

(Phosphine) does not form hydrogen bonds (with water)

25
Q

state the type of intermolecular force of attraction
HF CH3Cl Ne

A

HF has hydrogen bonding forces of attraction between molecules
CH3Cl has permanent dipole-dipole forces of attraction between molecules
Neon has induced dipole dipole forces of attraction between atoms

26
Q

This question is about the periodicity of the Period 3 elements.

(a) State and explain the general trend in first ionisation energy across Period 3

A

General increase

If not increase then CE
Greater nuclear charge / more protons

Same shielding / electrons added to same shell

Stronger attraction (from nucleus) for outer electron(s)

27
Q

Give one example of an element which deviates from the general trend in first ionisation energy across Period 3.

Explain why this deviation occurs.

(3

A

Aluminium / Al
(Outer) electron in (3)p orbital / sub-shell (level)

(3p) higher in energy

Allow more shielded or weaker nuclear attraction

OR

Sulfur / S (lower than P)
(Outer) electrons in (3)p orbital begin to pair
Repel

28
Q

In terms of structure and bonding, explain why calcium has a higher melting point than potassium.

A

M1 Both have giant metallic lattice structure, with metallic bonds / electrostatic attractions between positive metal ions and delocalized electrons

M2 Calcium has greater ionic charge (+2) than potassium (+1) and more delocalized electrons

M3 Therefore, stronger metallic bonds and more energy needed to break them

29
Q

SbCl3 molecules are polar. Explain why

A

There is a difference in electronegativities (between Sb and Cl)
OR
(Sb-Cl) bonds are polar OR have a dipole

M2 The molecule is not symmetrical
AND
The dipoles do not cancel

30
Q

Nickel is a metal with a high melting point.
(ii)
Explain, in terms of its structure and bonding, why nickel has a high melting point.

A

Contains positive (metal) ions and delocalised / mobile / free electrons
Strong attraction between them or strong metallic bonds

31
Q

Explain why nickel is ductile (can be stretched into wires).

A

Layers / planes / sheets of ions can slide over one another

32
Q

Suggest why the electronegativity of the elements increases from lithium to fluorine.

A

More protons / bigger nuclear charge
Same or similar shielding / electrons in the same shell or principal energy level / atoms get smaller

33
Q

Deduce why the bonding in nitrogen oxide is covalent rather than ionic.

A

Small electronegativity difference

34
Q

There are two lone pairs of electrons on the oxygen atom in a molecule of oxygen difluoride (OF2).
Explain how the lone pairs of electrons on the oxygen atom influence the bond angle in oxygen difluoride.

A

Lone pairs repel more than bond pairs
bond angle will be lower (than regular tetrahedral angle)

35
Q

Suggest why methaneselenol (CH3SeH) has a higher boiling point than methanethiol (CH3SH)

A

Methaneselenol is a) bigger molecule / larger Mr / larger no of electrons / Se bigger atom

With stronger/more vdw forces between molecules

36
Q

Explain, in terms of structure and bonding, why the melting point of sodium fluoride is
high.
[2 marks]

A

(Electrostatic) forces of attraction between oppositely charged
ions/Na+
and F–
Lots of energy needed to overcome/break forces

37
Q

Na+ and F– have the same electron configuration.
Explain why a fluoride ion is larger than a sodium ion.
[2 marks

A

Fluoride ion has (two) fewer protons/lower nuclear charge
Weaker attraction between nucleus and (outer) electrons
JUST DO NASA

38
Q

BrF4– ions are also formed when potassium fluoride dissolves in liquid BrF3 to form
KBrF4
Explain, in terms of bonding, why KBrF4 has a high melting point. 3

A

Ionic or (forces of) attraction between ions / bonds between ions

If molecules, IMF, metallic, CE =0

If covalent bonds mentioned, 0/3, unless specified within the BrF4– ion and not broken

Ignore atoms

Strong (electrostatic) attraction / strong bonds / lots of energy needed to break bonds
Between K+ and BrF4– ions/oppositely charged ions / + and – ions
If ions mentioned they must be correct
Strong bonds between + and – ions =3/3
(must know ionic is > covalent Q implying why its high so talk about ionic)

39
Q

Titanium is also a strong material that has a high melting point. It has a structure similar to that of magnesium.
Titanium can be hammered into objects with different shapes that have similar
strengths. 1
b)Suggest why these objects with different shapes have similar strengths.1

A

Layers of atoms/ions slide (over one another)
DUCTILE like gold (jewellery in many shapes)

b)(Strong) (metallic) bonding re-formed / same (metallic) bonding / retains same (crystal) structure / same bond strength / same attraction between protons and delocalised electrons as before being hammered or

40
Q

Explain why calcium has a higher melting point than strontium.

A

Calcium has a higher melting point than strontium, because
Ignore general Group 2 statements.
Correct reference to size of cations / proximity of electrons
M1 (For Ca) delocalised electrons closer to cations / positive ions / atoms /
nucleus
OR cations / positive ions / atoms are smaller
OR cation / positive ion / atom or it has fewer (electron) shells / levels
Penalise M1 if either of Ca or Sr is said to have more or less
delocalised electrons OR the same nuclear charge.
Ignore reference to shielding.
Relative strength of metallic bonding
M2 (Ca) has stronger
attraction between the cations / positive ions / atoms /
nucleus
OR
and the delocalised electrons
stronger metallic
bonding

CHEM (20 decks)

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