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Physical Chemistry > Bonding 1 > Flashcards

Bonding 1 Flashcards

1
Q

What are the three types of strong chemical bond

A

Ionic, covalent and metallic

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2
Q

Describe ionic bonding

A
  • Ionic bonding occurs between metals and non-metals.
  • Metal atoms lose electrons to form positive ions and non-metal atoms gain these electrons to form negative ions.
  • Both ions have a full (stable) electron structure.
  • Electrostatic forces cause the oppositely charged ions to be attracted to each other.
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3
Q

What structure to ionic compounds always exist in

A

A lattice

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4
Q

Why are ionic compounds always solids at room temperature

A

They have giant structures and therefore high melting temperatures because in order to melt an ionic compound, energy must be supplied to break up the lattice of ions.

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5
Q

Why do ionic compounds conduct electricity when molten or aqueous but not when they are a solid

A

This is because the ions that conduct electricity are free to move in a liquid but cannot move in a solid.

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6
Q

Why are ionic compounds brittle (they shatter easily when given a sharp blow)

A

Ionic compounds form a lattice of alternating positive and negative ions.
Therefore a blow in certain direction could move the ions and produce contact between ions with like charges, causing the compound to shatter.

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7
Q

List the three key properties of ionically bonded compounds

A
  • They are solids at room temperature
  • They conduct electricity when molten or aqueous (but not solid)
  • They are brittle.
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8
Q

What is a covalent bond

A

A shared pair of electrons

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9
Q

Describe covalent bonding

A
  • Covalent bonding happens between two non-metals.
  • the atoms share some of their outer electrons in covalent bonds so that each atom has a stable noble gas electron arrangement
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10
Q

How does covalent bonding (sharing electrons) hold atoms together

A

Atoms with covalent bonds are held together by the electrostatic attraction between the nuclei and the shared electrons.

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11
Q

What is a double bond

A

A bond in which four electrons are shared

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12
Q

List and describe the properties of substances with molecular structures

A
  • Substances composed of molecules can be solids, liquids and gases with low melting temperatures because the covalent bonds are only between the atoms within the molecules and there is only weak attraction between the molecules so the molecules do not need much energy to move apart from each other.
  • They are poor conductors of electricity because the molecules are equal overall. This means that there are no charged particles to carry the current.
  • If they dissolve in water, and remain as molecules, the solutions do not conduct water. Again, this is because there are no charged particles.
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13
Q

What is co-ordinate/dative covalent bonding

A

A covalent bond where one atom provides both of the electrons

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14
Q

What happens in co-ordinate/dative covalent bonding

A
  • the atom that accepts the electron pair is an atom that does not have a filled outer main level of electrons- it is electron deficient.
  • the atom that donates the electrons has a lone pair of electrons.
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15
Q

How are co-ordinate/dative covalent bonds represented

A

Instead of the straight line that represents normal covalent bonds, dative bonds are represented by a straight arrow that points towards the atom that is accepting the electron pair.

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16
Q

How does the length and strength of co-ordinate/dative covalent bonds compare with that of normal covalent bonds

A

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

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17
Q

Describe metallic bonding

A

Metallic bonding consists of a lattice fog positively charged metal ions surrounded by a ‘sea’ of delocalised electrons that are free to move and not associated with any one atom. The number of delocalised electrons depends on how many electrons have been lost by each metal atom. The metallic bonding spreads throughout so metals have giant structures.

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18
Q

List the four key properties of metals

A
  • Good conductors of electricity and heat
  • Strong
  • malleable and ductile
  • High melting point
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19
Q

Explain why metals are good conductors of electricity

A

The delocalised electrons can move throughout the structure. An electron from the negative terminal of the electrical supply joins the electron sea and at the same time, a different electron leaves the wire at the positive terminal.

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20
Q

Explain why metals are good conductors of heat

A

They have high thermal conductivity. This is partly due to the electron sea and energy is also spread by increasingly vigorous vibrations of the closely packed ions.

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21
Q

Describe and explain what factors affect the strength of a metal

A

1) The charge on the ion- the greater the charge on the ion, the greater the number of delocalised electrons and the stronger the electrostatic attraction between the positive ions and electrons.
2) The size of the ion- the smaller the ion, the closer the electrons are to the positive nucleus and the stronger the bond.

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22
Q

Explain why metals tend to be strong

A

The delocalised electrons extend throughout the solid so there are no bonds to break.

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23
Q

Why do metals generally have high melting and boiling points

A

Because they have giant structures. There is strong attraction between metal ions and the delocalised sea of electrons. This makes the atoms difficult to separate.

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24
Q

Define electronegativity

A

Electronegativity is the power of an atom to attract the electron density in a covalent bond towards itself.

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25
Q

Define electron density

A

Electron density is a representation of the probability of finding an electron in a specific location around an atom or molecule.

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26
Q

What is the Pauling scale

A

A measure of electronegativity that runs from 0 to 4. The greater the number, the more electronegative the atom.

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27
Q

What three factors does the electronegativity of an atom depend on

A

1) The nuclear charge
2) The distance between the nucleus and the outer shell electrons
3) The shielding of nuclear charge by electrons in inner shells

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28
Q

How does the size of an atom affect its electronegativity

A

The smaller the atom, the closer the nucleus is to the shared outer main level electrons and the greater its electronegativity.

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29
Q

How does the nuclear charge (for a given shielding affect) effect an atoms electronegativity

A

The larger the nuclear charge, the greater the electronegativity.

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30
Q

What is the trend in the electronegativity of atoms up/down a group in the periodic table

A
  • Going up a group in the periodic table, the electronegativity increases because the atoms get smaller and there is less shielding by electrons in inner shells.
  • Going down a group in the periodic table, the electronegativity decreases because the atoms get bigger and so there is more shielding by electrons in inner shells.
31
Q

What is the trend in electronegativity going across a period on the periodic table

A

Going across a period in the periodic table, the electronegativity increases. The nuclear charge increases, the number of inner main levels remain the same and the atoms become smaller.

32
Q

Where are the most electronegative atoms found on the periodic table

A

The top right corner

33
Q

What are the most electronegative atoms

A

Fluorine, oxygen and nitrogen followed by chlorine.

34
Q

What type of covalent bonds are completely non-polar

A

Covalent bonds between atoms of the same element

35
Q

What are the three types of intermolecular forces (in order from weakest to strongest)

A
  • Van der Waals
  • Dipole-Dipole
  • Hydrogen bonding
36
Q

How does the electronegativities of the atoms affect the polarity of a covalent bond

A

The greater the difference in the electronegativities of the two atoms in the covalent bond, the more polar the bond is

37
Q

What is the dipole moment of a molecule

A

The total affect that the polarity of all the bonds has on the molecule

38
Q

What molecules do Dipole-Dipole forces act between

A

Molecules that have permanent dipoles

39
Q

Describe what Van der waals forces are

A

Van der Waals forces describe the weak electrostatic attractions between all atoms and molecules

40
Q

List the key features of Van der Waals forces

A

1) Van der Waals act between all atoms and molecules at all times
2) They are in addition to any other intermolecular forces
3) The dipole is caused by the changing position of the electron cloud, so the more electrons there are, the larger the instantaneous dipole will be. Therefore the size of the Van der Waals forces increases with the number of electrons present.

41
Q

What atoms or molecules produce the strongest Van der Waals forces

A

Atoms or molecules with a large atomic or molecular masses as these contain a higher number of electrons.

42
Q

What two key things does the fact that the strength of the Van der Waals increases with the atom/molecules size explain

A
  • the boiling points of the noble gases increase as the atomic numbers of the noble gases increase.
  • the boiling points of hydrocarbons increase with increased chain length.
43
Q

List the atoms that are electronegative enough to form a hydrogen bond

A

Oxygen (O)
Nitrogen (N)
Fluorine (F)

44
Q

What conditions must be satisfied for hydrogen bonding to occur

A

For hydrogen bonding to occur, there must be a very electronegative atom with a lone pair of electrons covalently bonded to a hydrogen atom.

45
Q

How strong are hydrogen bonds in relation to other intermolecular forces/bonds

A

Hydrogen bonds are stronger than dipole-dipole forces but weaker than a covalent bond

46
Q

How do the boiling points of hydrides of elements in group 4,5,6 and 7 show the effects of hydrogen bonding

A
  • the boiling points of water (H2O), hydrogen fluoride (HF) and ammonia (NH3) are all higher than the hydrides of other elements in their group, but you would expect them to be lower if only Van Der Waals forces were operating.
  • This is because hydrogen bonding is present between molecules in each of these compounds and these stronger intermolecular forces of attraction make the molecules more difficult to separate.
47
Q

Describe why the density of ice is lower than the density of water (and water expands when frozen)

A
  • When water is in its liquid state the hydrogen bonds break and reform easily as the molecules are free to move.
  • When the water freezes, the molecules are not longer free to move and the hydrogen bonds hold them in fixed positions.
  • This holds the molecules in a 3D structure where they are less closely packed than in liquid water.
  • This means water expands when frozen and makes ice less dense than water.
48
Q

What does electron pair repulsion theory dictate

A
  • Each pair of electrons around an atom will repel all other electron pairs
  • The pairs of electrons will therefore take up positions as far apart as possible to minimise repulsion.
49
Q

What is the first step in finding the shape of a molecule

A

Draw a dot and cross diagram to work out the number of electron pairs present around the central atom.

50
Q

If there are two pairs of shared electrons around the central atom in a molecule, what is its shape and bond angle

A

Shape: Linear
Bond angle: 180 degrees
Example: Beryllium chloride BeCl2

51
Q

If there are three pairs of shared electrons around the central atom, what will the molecules shape and bond angle be

A

Shape: Trigonal planar
Bond angle: 120 degrees
Example: Boron triflouride BF3

52
Q

If there are four shared pairs of electrons around the central atom, what will the molecules shape and bond angle be

A

Shape: Tetrahedral
Bond angles: 109.5
Example: Methane and ammonium ion

53
Q

If there are five shared pairs of electrons around the central atom, what will the molecules shape and bond angles be

A

Shape: Trigonal pyramidal
Example: Phosphorus pentachloride PCl5

54
Q

If there are six shared pairs of electrons around the central atom,what will the molecules shape and bond angles be

A

Shape: Octahedral
Bond angles: 90 degrees
Example: Sulphur hexafluoride SF6

55
Q

If there are four electron pairs around the central atom but three are bonding pairs and one is a lone pair,what will the shape of the molecule and its bond angles be

A

Shape: Triangular pyramid
Bond angles: 107 degrees
Example: Ammonia NH3

56
Q

What is the approximate rule for the affect of lone pairs on bond angles

A

Lone pairs decrease the bond angles by roughly 2 degrees

57
Q

Explain why the bond angles in ammonia are 107 degrees when in a tetrahedron they are normally 109.5

A
  • The three bonding pairs of electrons are attracted by both the nitrogen and the hydrogen nuclei
  • The lone pair of electrons is only attracted by the nitrogen nucleus and is therefore pulled closer to it than the shared pairs
  • So repulsion between a lone pair of electrons and a bonding pair is greater than between two bonding pairs.
  • This effect squeezes the hydrogen atoms together, reducing the H-N-H bond angles to roughly 107 degrees.
58
Q

If there are four electron pairs around the central atom when two are bonding pairs and two are lone pairs, what is the molecules shape and bond angle

A

Shape: V-shaped
Bond angle: 104.5
Example: Water H20

59
Q

Explain why the shape of water is V-shaped

A
  • There are four electron pairs around the oxygen atom so the shape is based on a tetrahedron
  • However two of the ‘arms’ of the tetrahedron are lone pairs that are not part of a bond.
  • This results in a V-shaped or angular molecule.
60
Q

If there are six bonding pairs of electrons in a molecules and four are bonding pairs, two are lone pairs, what will the molecules shape and bond angle be

A

Shape: Square planar
Example: Chlorine tetraflouride ion ClF4-

61
Q

Explain why the shape of the chlorine tetrafluoride ion is square planar

A
  • There are four bonding pairs of electrons and two lone pairs
  • one of the lone pairs contains an electron that has been donated to it so the charge on the ion is -1
  • As there are six electron pair, the shape is based on an octahedron.
  • However, two ‘arms’ are not part of a covalent bond.
  • As lone pairs repel the most, they adopt the position furthest apart.
  • This leaves a square shaped ion called square planar.
62
Q

List which electron pairs repel each other the most in ascending order

A

Bonding pair- Bonding pair
Lone pair- Bonding pair
Lone pair- Lone pair

63
Q

Describe what happens when you heat a solid

A

1) When you first heat a solid, it makes the particles vibrate more (due to having more kinetic energy) around a fixed position. This slightly increases the average distance between particles and so the solid expands.
2) Energy (enthalpy change of melting) is supplied which weakens the forces between particles and allows them to move more freely in liquid form.
3) While a solid is melting, the temperature does not change because the heat energy provided is absorbed as the forces between particles are weakened.

64
Q

What is enthalpy

A

Heat energy change measured under constant pressure.

65
Q

Describe what happens when you heat a liquid

A
  • When a liquid is heated, the particles move further apart and so the liquid expands
  • In order for a liquid to turn into a gas, all of the intermolecular forces between particles must be broken
  • This requires the enthalpy change of vaporisation.
66
Q

Describe what happens when you heat a gas

A

When a gas is heated, the particles gain kinetic energy and move faster. They get much further apart and so gases expand a lot when heated.

67
Q

What are the four basic crystal types

A

1) Ionic
2) Metallic
3) Molecular
4) Macromolecular

68
Q

Describe the properties of ionic crystals/compounds

A
  • Ionic compounds have strong electrostatic attraction between oppositely charged ions
  • This strong electrostatic attraction means that ionic compounds have high melting points.
  • Ionic compounds only conduct electricity when molten or aqueous as this is when the ions are free to move.
69
Q

Describe the structure and properties of metallic crystals/compounds

A
  • Metals exist as a lattice of positive ions embedded in a delocalised sea of electrons.
  • They have high melting points due to the strong metallic bonds.
  • As the delocalised electrons are free to move, metals conduct electricity in all states.
70
Q

Describe the structure and properties of molecular crystals

A
  • Molecular crystals consist of molecules held in regular array by intermolecular forces.
  • Covalent bonds within the molecules hold the atoms together but do not act between the molecules.
  • Intermolecular forces are much weaker than metallic or ionic bonds so molecular crystals have low melting points and enthalpies of melting.
71
Q

Describe the structure and properties of a macromolecular compound

A
  • A macromolecular structure is one in which large numbers of atoms are linked in a regular three-dimensional arrangement by covalent bonds that extend throughout the compound.
  • These strong covalent bonds mean macromolecular crystals have high melting points.
72
Q

Give some examples of macromolecular compounds

A

Diamond, graphite, silicon dioxide.

73
Q

Describe the structure and properties of diamond

A
  • Diamond is a macromolecular structure
  • Each carbon atom forms four covalent bonds with other carbon atoms.
  • These four electron pairs repel each other (following the rules of electron pair repulsion theory) and form a tetrahedral shape with bond angles of 109.5 degrees.
  • The covalent bonds within the lattice are strong and mean that diamond has these properties:
    1) is a very hard material
    2) Has an incredibly high melting point
    3) does not conduct electricity because there are no free charged particles to carry charge.
74
Q

Describe the structure and properties of graphite

A
  • Graphite has both strong covalent bonding and weaker Van Der Waals forces.
  • Each carbon atom forms three single covalent bonds to other carbon atoms.
  • As predicted by electron pair repulsion theory, graphite forms a Trigonal planar structure, with bond angles of 120 degrees.
  • This leaves each carbon atom with a ‘spare’ electron in a p-orbital that is not part of the three single covalent bonds.
  • These ‘spare’ electrons merge above and below the plane of the carbon atoms in each layer. These electrons are delocalised, but in two dimensions only.
  • These delocalised electrons mean graphite can conduct electricity. They can travel freely through the material, but graphite will only conduct along the hexagonal planes, not at right angles to them.
  • There is no covalent bonding between the layers of carbon atoms- just weak Van Der Waals forces. This means that the layers can slide over one another making graphite soft and flaky.

Physical Chemistry (10 decks)

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