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Chemistry > Bonding and structure > Flashcards

Bonding and structure Flashcards

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1
Q

What is ionic bonding?

A

The strong electrostatic attraction between oppositely charged ions.

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2
Q

How does the charge of an ion effect the strength of an ionic bond?

A

Ions with a greater charge will have a greater attraction to other ions resulting in stronger forces of electrostatic attraction.

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3
Q

How does the ionic radius effect the strength of an ionic bond?

A

ions with greater atomic radius will have a weaker attraction to the oppositely charged ion because the attractive forces act over a greater distance.

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4
Q

What provides evidence that ions exist?

A

Physical properties of ionic compounds and the migration of ions.

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5
Q

What is a covalent bond?

A

A strong electrostatic attraction between two nuclei and the shared pair of electrons between them.

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6
Q

How does bond length effect bond strength of a covalent bond?

A

Shorter bonds are usually stronger as the atoms are held closer together so the forces of attraction are greater. Double and triple bonds are generally shorter than single bonds.

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7
Q

What intermolecular force do all simple covalent structures experience?

A

London forces

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8
Q

What determines the shape of a simple molecule?

A

The shape is determined by the number of electron pairs around the central atom. The bonding pair of electrons separate out to maximise the separation. So the largest possible bond angle exists between covalent bonds.

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9
Q

What does lone pair repulsion do and why?

A

Reduce bond angle by approximately 2.5 degrees per pair this is due the fact lone pair of electrons are held more closely to the central atom and so exert a greater repelling effect than bonding pairs.

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10
Q

State the following for linear molecules:
- number of e- bonding pairs
- number of e- lone pairs
-bond angle
-key examples

A
  • e- bonding pairs = 2
  • e- lone pairs = 0
  • bond angle = 180 degrees
  • examples - becl2, co2
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11
Q

State the following for v - shaped molecules:
- number of e- bonding pairs
- number of e- lone pairs
-bond angle
-key examples

A
  • e- bonding pairs = 2
  • e- lone pairs = 2
  • bond angle = 104.5 degrees
  • examples - H2O
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12
Q

State the following for Triangular pyramid:
- number of e- bonding pairs
- number of e- lone pairs
-bond angle
-key examples

A
  • e- bonding pairs = 3
  • e- lone pairs = 1
  • bond angle = 107 degrees
  • examples - NH3
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13
Q

State the following for Tetrahedral:
- number of e- bonding pairs
- number of e- lone pairs
-bond angle
-key examples

A
  • e- bonding pairs = 4
  • e- lone pairs = 0
  • bond angle = 109.5 degrees
  • examples - CH3
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14
Q

State the following for Trigonal Bipyramidal:
- number of e- bonding pairs
- number of e- lone pairs
-bond angle
-key examples

A
  • e- bonding pairs = 5
  • e- lone pairs = 0
  • bond angle = 3 -120 and 2- 90 degrees
  • examples - PCl5
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15
Q

State the following for Octahedral:
- number of e- bonding pairs
- number of e- lone pairs
-bond angle
-key examples

A
  • e- bonding pairs = 6
  • e- lone pairs = 0
  • bond angle = 90 degrees
  • examples - SF6
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16
Q

What is electronegativity?

A

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

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17
Q

What happens to electronegativity along a period?Why?

A

Increases as atomic radius decreases and the proton number increases.

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18
Q

What happens to electronegativity down a group?

A

decreases as atomic radius increases and the shielding increases.

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19
Q

When will a polar covalent bond form?

A

If the difference in electronegativity is between 0.4 and 1.7,

20
Q

When will an ionic bond form?

A

If the difference in electronegativity is greater than 1.7.

21
Q

When do polar molecules arise?

A

When there is an overall polarity across the molecule. Polar molecules MUST have polar bonds but a molecule with polar bonds may not always form polar molecules.

22
Q

How do London forces work?

A
  • Uneven distribution of electrons due to their random movement
    -Results in an instantaneous dipole being set up in that molecule
    -This induces a dipole on another molecule
23
Q

What two factors affect the strength of London forces and why do they affect the strength?

A

-More electrons means stronger forces - Dipoles created will be greater.
-Shape - long thin molecules have larger instantaneous dipoles as can approach closer. Reducing the distance the force has to act over.

24
Q

When do permanent dipole - permanent dipole forces form?

A

When a polar molecule forms one atom is permanently more negative than the other meaning that a delta - and a delta + region forms. act in addition to London forces.

25
Q

What must molecules have to form hydrogen bond between one another?

A

-Either O,N or F atom will have lone pairs of electrons
-The O,N and F atoms attract electrons from small hydrogen atoms so strongly they leave exposed protons. These protons strongly attract lone pairs from neighbouring O,N and F atoms.

26
Q

What is a hydrogen bond?

A

A particular strong permanent dipole - permanent dipole attraction between the lone pair of electrons on a very electronegative atom (N,O or F) and a hydrogen atom directly covalently bonded to another very electronegative N,O or F atom

27
Q

Why does ice have a lower density than water?

A

Hydrogen bonds - hold the molecules in a rigid structure as ice with lots of air gaps. As ice begins to melt some hydrogen bonds break and the structure collapses in on itself.

28
Q

Why do all alcohols have much higher boiling points then alkanes with similar molecular mass values?

A

lone pair on oxygen able to form hydrogen bonds

29
Q

Why can water dissolve most ionic compounds?

A

Hydrogen bonding allows it to dissolve them by solvating the individual ions

30
Q

Why can water dissolve some alcohols?

A

Hydrogen bonding allows it to dissolve them by forming hydrogen bonds with their hydroxyl group

31
Q

What are water and alcohols poor at dissolving?

A

Polar molecules that can not from hydrogen bonds.

32
Q

When are non-aqueous solvents used?

A

Used with compounds which have the same type of intermolecular force.

33
Q

Explain the trend in boiling temperatures of the hydrogen halides?

A

-HF highest as forms hydrogen bonds
-increases from HCl to HI due to increased strength of London forces due to more electrons being present.

34
Q

What is metallic bonding?

A

Strong electrostatic attraction between metal ions and sea of delocalised electrons

35
Q

What two factors effect the strength of metallic bonding?

A
  • charge - greater charge on ion means greater attraction between delocalised electrons and ions
  • size of ion - greater atomic radius decreases attraction due to lower charge density
36
Q

What are giant lattices present in?

A
  • ionic solids
    -covalently bonded solids
    -solid metals
37
Q

Why do ionic structures have a high melting point?

A

Electrostatic attraction is strong and requires a lot of energy to overcome

38
Q

Why do ionic structures conduct electricity when dissolved or in molten?

A

-Ions separate and are no longer in the lattice
-ions are free to move and carry the charge

39
Q

Why are ionic structures usually brittle?

A

-When layers of alternating charges are distorted the like charges repel one another - breaking the lattice into fragments

40
Q

Why are metallic structures good conductors?

A

Delocalised electrons can move and carry the charge

41
Q

Why are metals malleable?

A

Layers of positive ions are able to slide over one another

42
Q

Why do metals have a high melting point?

A

Strong electrostatic force between ions and electrons

43
Q

Why are simple molecular structures poor conductors?

A

Contains no charged particles

44
Q

Why do simple molecular structures have a low melting point?

A

Intermolecular forces are very weak

45
Q

Why is diamond one of the strongest hardest materials?

A

Each carbon atoms is covalently bonded to four other carbon atoms forming a rigid tetrahedral structure

Chemistry (7 decks)

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