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Chemistry AS > Bonding > Flashcards

Bonding Flashcards

1
Q

What does ionic bonding occur between?

A

Metals and non-metals

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2
Q

In ionic bonding, where are the electrons transferred from and to?

A

Electrons from metals atoms to non-metals atoms

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3
Q

What is formed as a result of ionic bonding?

A

Positive and negative ions formed

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4
Q

In ionic bonding, what are the forces which attract the two ions?

A

Electrostatic forces

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5
Q

What structure do ionic compounds exist in?

A

Lattice

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6
Q

What state are ionic compounds in at room temperature?

A

Solid

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7
Q

What type of melting point do ionic compounds have? Why?

A
  • High melting point
  • Because in order to melt an ionic compound, energy must be supplied to break up the lattice of ions
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8
Q

Do ionic compound conduct electricity?

Why?

A
  • Ionic compounds only conduct electricity when molten or dissolved in water (aqueous), not when solid
  • This is because the ions that carry the current are free to move in the liquid start but are not free in the solid state
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9
Q

Are ionic compounds strong or weak?

Why?

A
  • Brittle and shatter easily when given a sharp blow
  • This is because they form a lattice of alternating positive and negative ions
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10
Q

What do covalent bonds form between?

A

A pair of non metal atoms

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11
Q

What are covalent bonds?

A

A covalent bond is shared pair of electrons (the atoms share some of their outer electrons so that each atom has a stable noble gas arrangement)

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12
Q

What are the charges of the atoms within a covalent bond?

A

Neutral (no electrons have been transferred from one atom to another

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13
Q

How is Methane gas bonded? Draw the bonding:

A

Covalently bonded

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14
Q

What holds covalent bonds together?

A

Electrostatic attraction between the nuclei and the shared electrons

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15
Q

How many electrons are shared in a double bond?

Example of this:

A

4 electrons are shared in a double bond

Oxygen, O2

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16
Q

Are covalently bonded molecules strongly or weakly attracted to each other?

A

Weakly attracted to each other

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17
Q

Do substances composed of molecules have high or low melting point?
Why?

A
  • Low melting point
  • Because strong covalent bonds are only between atoms within the molecules
  • There is only weak attraction between the molecules so the molecules do not need much energy to move apart from each other
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18
Q

Are substances, composed of molecules, good or poor conductors?

Why?

A
  • Poor conductors of electricity
  • Because the molecules are neutral overall
  • No charged particles to carry the current
  • Doesn’t change when dissolved in water and remain as molecules as no charges particles
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19
Q

What is co-ordinate bonding/dative bonding?

A

Where one atom provides both the electrons

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20
Q

What are the features of an atom in a dative bond?

A
  • The atom accepts the electron pair is an atom that does not have a filled outer main level of electrons - the atom is electron deficient
  • The atom that is donating the electrons has a pair electrons that is not being used in a bond (called lone pair)
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21
Q

What does electron deficient mean?

A

Does not have a filled outer main level of electrons

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22
Q

How are co-ordinate/dative bonds represented?

A
  • With an arrow, pointing towards the atom that is accepting the electron pair
  • This only shows how the bond was made
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23
Q

Do co-ordinate bonds have same or different strength and length as normal covalent bonds between the same pair of atoms?

A

Same

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24
Q

Draw an example (of the dative bond in the ammonium ion)

A
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25
Q

How do you describe metallic bonding?

A

A lattice of positive ions existing in a ‘sea’ of delocalised outer electrons. The positive ions tend to repel one another and this is balanced by the electrostatic attraction of these positive ions for the negatively charged ‘sea’ of delocalised electrons.

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26
Q

What does the number of delocalised electrons depend on in metallic bonding?

A

Depends on how many electrons have been lost by each metal atom

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27
Q

Is metallic bonding a giant structure?

A

Yes as it spreads throughout

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28
Q

Why are metals good at conducting electricity?

A

The delocalised electrons can move throughout the structure

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29
Q

Why are metals good conductors of heat?

A

They have high thermal conductivities. The sea of electrons is partly responsible for this property, with energy also spread by increasingly vigorous vibrations of the closely packed ions.

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30
Q

What does the strength of a metal depend on?

A
  • The charge on the ion. The greater the charge, the greater the number of delocalised electrons and the stronger the electrostatic attraction between the positive ions and electrons
  • 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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31
Q

Why are metals strong?

A

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

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32
Q

Why are metals malleable and ductile?

A

After a small distortion, each metal ion is still in exactly the same environment as before, so the new shape is retained.

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33
Q

Why do metals have high melting and boiling points?

A

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

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34
Q

Definition of electronegativity

A

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

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35
Q

What is the Pauling scale used for?

Explain it

A

As a measure of electronegativity. It runs from 0 to 4. Greater the number, the more electronegative the atom.

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36
Q

Do noble gases have a number on the Pauling Scale?

Why?

A
  • No
  • They do not (in general) from covalent bonds
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37
Q

What does Electronegativity depend on?

A
  1. Nuclear charge (larger the nuclear charge, greater its electronegativity)
  2. Atomic radius (smaller the atom, greater its electronegativity)
  3. Shielding
38
Q

What is the trend in electronegativity?

A
  • Going up a group, electronegativity increases

(atoms get smaller, less shielding)

  • Going across a period the electronegativity increases

(nuclear charge increases, atoms get smaller)

39
Q

Where are the most electronegative atoms found on the periodic table?

A

Top right corner

40
Q

What is a polar bond?

A

A covalent bond between two atoms in which the electrons in the bond are unevenly distributed. This causes a slight charge difference, inducing a dipole in the molecule

41
Q

What is polarity?

A

The unequal sharing or the electrons between atoms that are bonded together covalently

42
Q

If the atoms are the same in a covalent bond (e.g., F2), how are the electrons shared and what happens with polarity and electronegativity?

A
  • Shared equally between atoms
  • Same electronegativity
  • Bond completely non-polar
43
Q

If the atoms are the different in a covalent bond, how are the electrons shared and what happens with polarity and electronegativity?

A
  • Electrons not shared equally between atoms
  • Electrons attracted more to atom with higher electronegativity
  • Electron cloud distorted towards the atom with higher electronegativity and that atom is delta -
  • Polar bonds
  • The greater the difference is electronegativity, the more polar is the covalent bond
44
Q

What are the three types of intermolecular forces?

A
  • van der Waals forces
  • Dipole-dipole forces
  • Hydrogen bonding
45
Q

Which type of intermolecular force can act between all atoms and molecules?

A

van der Waals forces

46
Q

Which is the weakest and strongest intermolecular force?

A

Weakest: van der Waals forces

dipole-dipole forces

Strongest: Hydrogen bonding

47
Q

Why would dipoles cancel?

A
  • If they are in opposite direction and of same magnitude (same atoms)
  • If the molecule is completely symmetric - therefore, the molecule is non polar
48
Q

What do dipole-dipole forces only act on?

A

Molecules that have permanent dipoles

49
Q

What are van der Waals forces?

A
  • Aka - induced dipole-dipole, dispersion, London forces
  • Exist between all molecules
  • Arise due to fluctuations of electron density within a non polar molecule
  • These fluctuations may temporarily cause an uneven electron distribution, producing an instantaneous dipole
  • This dipole can induce a dipole in another molecule, and so on
  • All atoms and molecules are made up of positive and negative charges even though overall neutral
  • Charges produce weak electrostatic attractions between all atoms and molecules (van der Waals forces)
  • van der Waals forces are in addition to any other intermolecular forces
50
Q

What do ‘van der Waals forces’ increase with?

A
  • With increase of electrons present
  • More electrons, the larger the instantaneous dipole will be
  • Atoms/molecules with large atomic/molecular masses produce stronger van der Waals forces than smaller masses
51
Q

Why do boiling points of noble gases increase as the atomic numbers of the noble gases increase?

A

Stronger van der Waals forces as there is a greater number of electrons. Therefore, the instantaneous dipole will be larger, with the greater number of electrons

52
Q

Why do boiling points of hydrocarbons increase with increased chain length?

A

More electrons, greater van der Waals forces.

53
Q

What does hydrogen bonding consist of?

A

A hydrogen atom ‘sandwiched’ between two very electronegative atoms.

54
Q

What conditions needed for hydrogen bonding to occur?

A
  • Need very electronegative atom with a lone pair of electrons covalently boded to a hydrogen atom
55
Q

Why is the intermolecular bonding (hydrogen bonding) stronger than the weak dipole-dipole attractions in between molecules of water?

A
  1. The oxygen atoms in water have lone pairs of electrons
  2. In water the hydrogen atoms are highly electron as oxygen is very electronegative and attracts the shared electrons in the bond towards it. The hydrogen atoms in water + charged and very small. These exposed protons have a very strong electric field because of their small size
56
Q

What must there be for a hydrogen bond to form?

A
  • A hydrogen atom that is bonded to a very electronegative atom. This will produce partial positive charge on the hydrogen atom
  • A very electronegative atom with a lone pair of electrons. These will be attracted to the partially charged hydrogen atom in another molecule and form the bond
57
Q

What atoms are electronegative enough to form hydrogen bond?

A
  • Oxygen, O
  • Nitrogen, N
  • Fluorine, F
58
Q

Which is stronger, covalent bonds or hydrogen bonds?

59
Q

Why is ice less dense than water?

A
  • Water in liquid state, hydrogen bonds break and reform easily as the molecules are moving about
  • Water as solid, hydrogen bonds hold the molecules in fixed positions
  • To fit into this structure, molecules are slightly less closely packed than as liquid, and therefore, less dense.
60
Q

What is the electron pair repulsion theory?

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
  • Electron pairs may be a shared pair or a lone pair
61
Q

What is the bond angle/shape of molecule with two pairs of electrons?

A

180 (linear)

62
Q

What is the bond angle/shape of molecule with three pairs of electrons?

A

120 (trigonal planar)

63
Q

What is the bond angle/shape of molecule with four pairs of electrons?

A

109.5 (tetrahedral)

64
Q

What is the bond angle/shape of molecule with five pairs of electrons?

A

90 and 120 (Trigonal bipyramid)

65
Q

What is the bond angle/shape of molecule with six pairs of electrons?

A

90 (octahedral)

66
Q

What do the wedges and dotted lines represent?

A

Wedges - coming out of page

Dotted line - going into page

67
Q

What is the bond angle/shape of molecule with 4 pairs of electrons (1 lone pair)?

A

107 (triangular pyramid)

68
Q

Why is the bond angle less when with a lone pair?

A
  • Bonding pairs of electrons are attracted towards the nitrogen nucleus and also the hydrogen nucleus
  • Lone pair is attracted only by the nitrogen and therefore pulled closer to it than the shared pairs
  • So repulsion between a lone pair of electrons and a bonding pair of electrons is greater than that between two bonding pairs
  • This effect squeezes the hydrogen atoms together, reducing angles
69
Q

What is the bond angle/shape of molecule with four pairs of electrons (two lone pairs)?

A

104.9 (V-shaped)

70
Q

What is the bond shape of molecule with 4 pairs of electrons (2 lone pairs)?

A

square planar

71
Q

What is the arrangement of particles in crystals?

A

Regular arrangement and help together by forces of attraction

71
Q

What is the arrangement of particles in crystals?

A

Regular arrangement and help together by forces of attraction

72
Q

What are the 4 types of crystal types?

A

Ionic, metallic, molecular and macromolecular

73
Q

Do ionic crystals have high or low melting points?
Why?

A
  • High melting point
  • Because of the strong electrostatic attraction which extend throughout the structure
  • Require a lot of energy to break in order for ions to move apart from each other
74
Q

Do metallic crystals have high or low melting points?
Why?

A
  • High melting point
  • Strong metallic bonds, attraction of positive to negative extends throughout the crystal
75
Q

Do molecular crystals have high or low melting points?
Why?

A
  • Low melting point
  • Low enthalpy of melting
  • Intermolecular forces much weaker than covalent, ionic or metallic bonds
76
Q

What are the properties of molecular crystals?

A
  • Soft and break easily
  • Low melting temperature and sublimes readily to form gaseous iodine molecules
  • Does not conduct electricity because there are no charge particles to carry charge
77
Q

Do macromolecular crystals have high or low melting points?
Why?

Example?

A
  • High melting temperature
  • Covalent bonds extend throughout the compound and have the typical property of a giant structure held together with strong bonds
  • Diamond and graphite
78
Q

What are the properties of diamond?

A
  • Very hard material
  • Very high melting temperature (over 3700K)
  • Does not conduct electricity because there are no free charged particles of carry charge
79
Q

What is the shape and bond angles of diamond?

A

Tetrahedron (109.5)

80
Q

How many bonds does graphite form?

A

3 single covalent bond to other carbon atoms

81
Q

How many bonds does diamond form with other carbons?

A

4 covalent bonds

82
Q

What shape/bond angle does graphite have?

A

Trigonal planar (120)

83
Q

What happens with the ‘spare’ electron in graphite bonding?

A
  • In p-orbital, not part of three single covalent bonds
  • The p-orbitals with the ‘spare’ electron marge above and below the plane of the carbon atoms in each layer
  • These electrons can move anywhere within the layer, they are delocalised
84
Q

Why can graphite conduct electricity?

A
  • The ‘spare’ delocalised electron can travel freely through the material
  • Graphite will only conduct along the hexagonal planes, not at right angles to them
85
Q

What forces hold the layer of carbon atoms in graphite?

A
  • van der Waals forces
86
Q

Why can layers of graphite slide over each other?

A

Weak intermolecular force of attraction means that the layer can slide across one another making graphite soft and flaky

87
Q

What are the properties of graphite?

A
  • Soft material
  • High melting point (giant structure)
  • Conducts electricity along planes of hexagons
88
Q

How are the C atoms arranged in buckminsterfullerene?

A

football-like shape

89
Q

What is the formula for buckminsterfullerene?

Chemistry AS (14 decks)

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