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Physical Chemistry > Physical 3: Bonding > Flashcards

Physical 3: Bonding Flashcards

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1
Q

What is ionic bonding?

A

The result of electrostatic attraction between oppositely charged positive metal ions and negative non-metal ions

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2
Q

How do electrons move in ionic bonding?

A

From the metal to the non-metal
Forming a positive metal ion
And a negative non-metal ion

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3
Q

What are electrostatic forces?

A

Forces of attraction between oppositely charged ions like metal ions and non-metal ions

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4
Q

Why are ionic compounds solid at room temperature?

A

Ionic compounds form a giant ionic lattice
Strong electrostatic forces of attraction between positive metal ions and negative non-metal ions in the lattice
Require lots of energy to overcome

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5
Q

Can ionic compounds conduct?

A

Not when solid
But can when molten or dissolved because the ions will be free to move and carry the charge throughout the lattice

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6
Q

Why are ionic compounds brittle?

A

If a force was applied, it could displace the ions in the lattice
So that ions with the same charge were in contact
Which would cause repulsion, causing the lattice to shatter

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7
Q

What is covalent bonding?

A

A shared pair of electrons between two atoms
Strong electrostatic attraction between the nuclei and shared electrons

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8
Q

What is a molecule?

A

A small group of covalently bonded atoms
For example, Cl2

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9
Q

Give examples of simple covalent molecules.

A

Hydrogen (H2)
Hydrogen Chloride (HCl)
Water (H2O)
Carbon Dioxide (O=C=O)
Ethene (CH2=CH2)

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10
Q

Why do covalent molecules have low melting points?

A

Because the strong covalent bonds are only within the molecule (between atoms)
The intermolecular forces of attraction between molecules are much weaker and therefore easier to overcome

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11
Q

Why are covalent molecules poor conductors of electricity?

A

Because simple covalent molecules have no overall charge
So have no charged particles to carry a current

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12
Q

What is co-ordinate bonding?

A

Where one atom provides a lone pair of electrons for a covalent bond

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13
Q

Describe the co-ordinate bonding in an ammonium ion.

A

H+ ions are electron deficient
So the lone pair on nitrogen in a molecule of ammonia are donated to the H+ ion
(Represented with an arrow from the lone pair to the electron-deficient species)

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14
Q

How do co-ordinate bonds compare to ordinary covalent bonds?

A

They have exactly the same strength and length as ordinary covalent bonds between the same pair of atoms

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15
Q

Why do atoms bond?

A

To obtain a full outer shell of electrons
(Noble gas arrangement)
To reach a more stable energy state

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16
Q

What is metallic bonding?

A

A lattice of positively charged metal ions in a sea of delocalised electrons
Positive ions tend to repel each other, balanced out by the strong electrostatic forces of attraction between ions and delocalised electrons

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17
Q

Why are metals good conductors?

A

Because delocalised electrons can move throughout the structure

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18
Q

Why are metals good thermal conductors?

A

Because delocalised electrons can rapidly transfer heat energy by colliding with neighbouring atoms and transferring kinetic energy

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19
Q

What are the two main factors that the strength of metallic bonding depend on?

A

1) Charge on the ion (higher charge = more delocalised electron, so stronger electrostatic forces of attraction between positive ions and sea of delocalised electrons)

2) Size of the ion (smaller ion = closer to electrons, so stronger electrostatic forces of attraction between positive ions and sea of delocalised electrons)

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20
Q

Why are giant metallic lattices so strong in general?

A

Because there are no individual bonds to break
There are electrostatic forces of attraction throughout the entire structure, which requires lots of energy to overcome

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21
Q

Why are metals malleable and ductile?

A

Because the ions are all the same size
So they are able to easily slide past each other in layers
Due to the mobility of delocalised electrons

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22
Q

What is electronegativity?

A

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

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23
Q

What three factors affect electronegativity?

A

1) Atomic radius (smaller = more electronegative)
2) Shielding (less shielding = more electronegative)
3) Nuclear charge (higher nuclear charge = more electronegative)

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24
Q

Describe the trend in electronegativity across the Periodic Table.

A

Electronegativity increases as you go from left to right
Because shielding stays the same
But nuclear charge increases
So forces of attraction between the outer electron and nucleus increase
And the power of the atom to attract electrons increases

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25
Q

Describe the trend in electronegativity down a group on the Periodic Table.

A

Electronegativity decreases as you go down a group
Because atomic radius increases as number of shells increases
And shielding increases as a result of this
So there is a weaker force of attraction between the outer electron and the positive nucleus
So the atom has a weaker attraction of electrons

26
Q

What does polarity of a bond depend on?

A

The difference in electronegativity between two covalently bonded atoms
Polarity will be symmetrical if two of the same atom are bonded because they have the same electronegativity
Hydrogen and fluorine have a large difference in electronegativity, so will form polar bonds

27
Q

What are the three main intermolecular forces?

A

1) Van der Waals
2) Dipole-dipole
3) Hydrogen bonding

28
Q

How do van der Waal forces of attraction form?

A

Distribution of charges change in an atom due to movement of the electron cloud
Can create an instantaneous dipole in one atom
Which induces dipoles in nearby atoms
Forming instantaneous dipole-dipole forces of attraction

29
Q

Why do boiling points of noble gases increase down the group?

A

Because down the group, the atomic number increases
So number of electrons increases
More electrons mean that van der Waal forces are stronger as the instantaneous dipoles are larger
More van der Waal forces require more energy to overcome

30
Q

What are dipole-dipole forces of attraction?

A

Molecules with polar bonds can have dipole moments
They only act between molecules with permanent dipoles caused by differences in electronegativity

31
Q

Why do symmetrical molecules not have dipole-dipole forces?

A

Because the polar bonds cancel out
There is no overall dipole in the molecule

32
Q

What are the conditions needed for hydrogen bonding?

A

Two very electronegative atoms with a lone pair of electrons bonded to hydrogen atoms
E.g. water, ammonia

33
Q

What three elements can be involved in hydrogen bonds?

A

Nitrogen
Oxygen
Fluorine

34
Q

Why is hydrogen bonding so strong (e.g. in water)?

A

1) Oxygen atoms in water have 2 lone pairs
2) Hydrogen atoms in water are highly electron deficient because oxygen is more electronegative so attracts bonding pairs of electrons
3) Making the hydrogen atoms positively charged
4) So they can attract the lone pairs on oxygen on other water molecules

35
Q

Describe and explain the shape of hydrogen bonding.

A

Linear
Because the pair of electrons in the X-H bond (where X is nitrogen, oxygen, or fluorine) repel those in the hydrogen bond

36
Q

Why is ice less dense than water?

A

When water is in its liquid state:
Molecules are free to move around
So hydrogen bonds easily break and reform

When water freezes:
Molecules are no longer free to move around
So hydrogen bonds remain in a fixed position
Molecules must spread apart further to allow 3D structure to form
Making ice less dense

37
Q

Why is the lower density of ice useful?

A

Allows ice to float on water
Ice can act as an insulating layer e.g. on a pond
To help fish to survive through winter

38
Q

Draw the hydrogen bonding between two water molecules.

A

1) Delta positives and delta minuses
2) Two lone pairs on each oxygen molecule
3) Dashed line between O and H
4) Linear bond

39
Q

What is electron pair repulsion theory?

A

Each pair of electrons around an atom repel all other pairs, so each pair will take up positions as far apart as possible

40
Q

How do lone pairs affect structure of a molecule?

A

Lone pairs of electrons repel more from bonding pairs than bonding pairs repel from themselves
So reduce bond angle by 2.5 degrees

41
Q

What would be the shape of a molecule with 2 bonding pairs?

A

Linear
Bond angle 180

42
Q

What would be the shape of a molecule with 3 bonding pairs?

A

Trigonal planar
Bond angle 120

43
Q

What would be the shape of a molecule with 4 bonding pairs?

A

Tetrahedral
Bond angle 109.5

44
Q

What would be the shape of a molecule with 5 bonding pairs?

A

Trigonal bipyramid
Bond angles 120 and 90

45
Q

What would be the shape of a molecule with 6 bonding pairs?

A

Octahedral
Bond angle 90

46
Q

What would be the shape of a molecule with 3 bonding pairs and 1 lone pair?

A

Trigonal pyramid
Bond angle 107

47
Q

What would be the shape of a molecule with 2 bonding pairs and 2 lone pairs?

A

Bent
Bond angle 104.5

48
Q

What would be the shape of a molecule with 2 bonding pairs and 3 lone pairs?

A

Linear
Bond angle 180

49
Q

What would be the shape of a molecule with 4 bonding pairs and 2 lone pairs?

A

Square planar
Bond angle 90
E.g. Xenon Tetrafluoride

50
Q

What would be the shape of a molecule with 5 bonding pairs and 1 lone pair?

A

Square pyramidal
Bond angle 90

51
Q

Give the order of the strength of repulsion in types of electron pair.

A

Weakest: Bonding pair - bonding pair
Second: Bonding pair - lone pair
Strongest: Lone pair - lone pair

52
Q

What are the steps needed to work out the shape of a molecule?

A

1) What group is the central atom in? (e.g. 5)
2) How many bonds does it make?
3) Does it have a charge?
4) Total these numbers
5) Divide by 2 to give number of bonding pairs
6) Any lone pairs will be left over from (2)
7) Deduce the shape and bond angle

53
Q

Describe the arrangement of particles in solids, liquids, and gases.

A

Solids
- Regular arrangement
- Closely packed together
- Vibrate about a point

Liquids
- Randomly arranged
- Closely packed together
- Rapidly moving around

Gases
- Randomly arranged
- Far apart
- Rapid movement, exert pressure

54
Q

Why do molecular crystals have lower melting points than metallic and ionic crystals?

A

Because the molecules are only held together by weak intermolecular forces like van der Waals, which require less energy to overcome than ionic bonds or metallic bonds

55
Q

Describe the structure and properties of iodine crystals.

A

Iodine has a large number of electrons
So forms many strong van der Waals forces of attractions
Which collectively provide strength to the crystal
Giving it a high melting point
However, van der Waals forces are more easily broken than covalent bonds
So iodine crystals break more easily
And have lower melting points

56
Q

Does iodine conduct electricity?

A

No
Because there are no charged particles to carry a charge

57
Q

What is an allotrope of carbon?

A

The different physical forms that carbon can exist in

58
Q

Describe the macromolecular structure of diamond.

A

Forms four covalent bonds
Tetrahedral structure (four bonds repel equally)
Bond angle 109.5
Very hard material
Very high melting point
Doesn’t conduct electricity due to no delocalised electrons

59
Q

Describe the macromolecular structure of graphite.

A

Three covalent bonds
One delocalised electron per atom
Bond angle 120
Conducts electricity because of delocalised electrons, which can travel freely through the planes
Very high melting point
Van der Waals forces hold 2D layers together, so layers can move over each other easily, which is why graphite is soft

60
Q

What is buckminsterfullerene?

A

Closed cages of carbon atoms
Or nanotubes

Physical Chemistry (12 decks)

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