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Chemistry A Level > Topic 2 > Flashcards

Topic 2 Flashcards

1
Q

What happens to metal atoms during ionic bonding?

A

Metal atoms lose electrons to form +ve ions.

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2
Q

What happens to non-metal atoms during ionic bonding?

A

Non-metal atoms gain electrons to form -ve ions.

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3
Q

What is the structure of ionic crystals?

A

Ionic crystals have the structure of giant lattices of ions.

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4
Q

How does the ionic radius of Mg change when it forms an ion?

A

Mg goes from 1s² 2s² 2p⁶ 3s² to Mg²⁺ 1s² 2s² 2p⁶.

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5
Q

How does the ionic radius of O change when it forms an ion?

A

O goes from 1s² 2s² 2p⁴ to O²⁻ 1s² 2s² 2p⁶.

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6
Q

What is ionic bonding?

A

Ionic bonding is the strong electrostatic force of attraction between oppositely charged ions formed by electron transfer.

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7
Q

What factors affect the strength of ionic bonding?

A

Ionic bonding is stronger and the melting points higher when the ions are smaller and/or have higher charges.

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8
Q

Why does MgO have a higher melting point than NaCl?

A

MgO has smaller ions (Mg²⁺ & O²⁻) with higher charges than those in NaCl (Na⁺ & Cl⁻).

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9
Q

Why are positive ions smaller than their atoms?

A

Positive ions are smaller because they have one less shell of electrons and a greater net force on remaining electrons.

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10
Q

Why are negative ions larger than their corresponding atoms?

A

Negative ions are larger because they have more electrons than the corresponding atom, leading to less attraction per electron.

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11
Q

What is the electronic structure of N³⁻, O²⁻, F⁻, Na⁺, Mg²⁺, and Al³⁺?

A

N³⁻, O²⁻, F⁻, Na⁺, Mg²⁺, and Al³⁺ all have the same electronic structure of the noble gas Ne.

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12
Q

What happens to ionic radii within a group in the periodic table?

A

The size of the ionic radii increases going down the group due to more shells of electrons.

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13
Q

What shows the likelihood of finding electrons in a region?

A

Maps that display contours of equal electron density.

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14
Q

How are the ions arranged in NaCl?

A

The ions are arranged in a regular pattern.

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15
Q

Which ion is larger in NaCl?

A

The chloride ions are larger than the sodium ions.

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16
Q

What are the physical properties of ionic compounds?

A

They have high melting points due to strong attractive forces between the ions.

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17
Q

What happens to ionic compounds when solid?

A

They are non-conductors of electricity because ions are held tightly and cannot move.

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18
Q

What happens to ionic compounds when in solution or molten?

A

They conduct electricity because ions are free to move.

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19
Q

What is a characteristic of ionic compounds regarding brittleness?

A

Ionic compounds are brittle and easy to cleave apart.

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20
Q

What color do Cu²⁺ ions migrate to?

A

They migrate to the negative electrode and appear blue.

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21
Q

What color do CrO₄²⁻ ions migrate to?

A

They migrate to the positive electrode and appear yellow.

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22
Q

What does the electron density map indicate about ions?

A

The ions are discrete or separate, as the electron density falls to zero between them.

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23
Q

What limitation does the electron density map have?

A

It does not show the edge of the ion, making it difficult to measure the radius.

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24
Q

What happens when potassium manganate solution is placed on moist filter paper?

A

The purple color of the MnO₄⁻ ion migrates to the positive electrode after ten minutes.

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25
Q

What is a covalent bond?

A

A covalent bond is strong and is caused by the electrostatic attraction between the bonding shared pair of electrons and the two nuclei.

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26
Q

How is the strength of a covalent bond demonstrated?

A

The strength of a covalent bond can be demonstrated by the high melting points of giant atomic structures like diamond and graphite.

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27
Q

Why do diamond and graphite have high melting points?

A

They have high melting points because they contain many strong covalent bonds in a macromolecular structure.

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28
Q

What does it take to break strong covalent bonds?

A

It takes a lot of energy to break the many strong bonds.

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29
Q

What do X-ray diffractions for the hydrogen molecule show?

A

X-ray diffractions for the hydrogen molecule show a high concentration of negative charge between H nuclei.

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30
Q

What is the effect of negative charge in covalent bonding?

A

The negative charge is strongly attracted by both nuclei, so attractive interactions exceed repulsive ones.

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31
Q

What is significant about electron density in covalent compounds?

A

In a covalent compound, there is significant electron density between the atoms.

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32
Q

How do multiple bonds affect bond strength and length?

A

Nuclei joined by multiple (i.e., double and triple) bonds have a greater electron density between them.

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33
Q

What results from greater electron density in multiple bonds?

A

This causes a greater force of attraction between the nuclei and the electrons between them, resulting in a shorter bond length and greater bond strength.

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34
Q

What is a dative covalent bond?

A

A dative covalent bond forms when the shared pair of electrons in the covalent bond come from only one of the bonding atoms.

A dative covalent bond is also called coordinate bonding.

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35
Q

Can you provide common examples of dative covalent bonds?

A

Common examples include NH₄⁺, H₂O, and NH₃BF₃.

These examples should be drawable.

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36
Q

What is the shape of the ammonium ion (NH₄⁺)?

A

The shape of the ammonium ion (NH₄⁺) is tetrahedral.

The dative covalent bond acts like an ordinary covalent bond when considering shape.

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37
Q

What is the direction of the arrow in a dative covalent bond?

A

The direction of the arrow goes from the atom that is providing the lone pair to the atom that is deficient.

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38
Q

How do two aluminium chloride (AlCl₃) molecules interact?

A

Two aluminium chloride (AlCl₃) molecules join together through two dative bonds to form the dimer.

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39
Q

What is electronegativity?

A

Electronegativity is the relative tendency of an atom in a covalent bond in a molecule to attract electrons to itself.

Measured on the Pauling scale (ranges from 0 to 4).

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40
Q

How does electronegativity change across a period?

A

Electronegativity increases across a period as the number of protons increases and the atomic radius decreases.

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41
Q

How does electronegativity change down a group?

A

Electronegativity decreases down a group due to increased distance between the nucleus and outer electrons and increased shielding of inner shell electrons.

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42
Q

What are the extremes of bonding types?

A

Ionic and covalent bonding are the extremes of a continuum of bonding type.

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43
Q

What determines the type of bond in a compound?

A

Differences in electronegativity between elements can determine where a compound lies on the bonding scale.

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44
Q

What type of bond forms with similar electronegativities?

A

A compound containing elements of similar electronegativity will be purely covalent.

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45
Q

What type of bond forms with very different electronegativities?

A

A compound containing elements with a very large electronegativity difference (> 1.7) will be ionic.

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46
Q

What is a polar covalent bond?

A

A polar covalent bond forms when the elements in the bond have different electronegativities (around 0.3 to 1.7).

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47
Q

What happens in a polar covalent bond?

A

A polar covalent bond has an unequal distribution of electrons, producing a charge separation (dipole) with + and - ends.

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48
Q

What characterizes symmetric molecules?

A

A symmetric molecule (all bonds identical and no lone pairs) will not be polar even if individual bonds within the molecule are polar.

49
Q

Which atoms are the most electronegative?

A

F, O, N, and Cl are the most electronegative atoms.

50
Q

What is the most electronegative element?

A

The most electronegative element is fluorine, given a value of 4.0.

51
Q

What happens to dipoles in a symmetric molecule?

A

The individual dipoles on the bonds ‘cancel out’ due to the symmetrical shape of the molecule, resulting in no net dipole moment.

52
Q

Give an example of a non-polar molecule.

A

CCl4 is a non-polar molecule, whereas CHCl3 is polar.

53
Q

In a polar compound, which element is the δ- end?

A

The element with the larger electronegativity in a polar compound will be the δ- end.

54
Q

Is CO2 a polar or non-polar molecule?

A

CO2 is a symmetrical molecule and is a non-polar molecule.

55
Q

What happens when a charged rod is brought close to a polar liquid?

A

The jet of liquid will be attracted to the electrostatic force of the rod, causing the dipoles in the polar molecules to align.

The negative ends will be attracted to the positive rod (or vice versa).

56
Q

What occurs with non-polar liquids when a charged rod is brought close?

A

Non-polar liquids will not be deflected or attracted to the charged rod.

57
Q

What are London forces?

A

London forces occur between all molecular substances and noble gases, but not in ionic substances.

They are also called instantaneous, induced dipole-dipole interactions.

58
Q

How do temporary dipoles form in molecules?

A

Electrons in a molecule move constantly and randomly, causing fluctuations in electron density and the formation of small temporary dipoles.

59
Q

What are induced dipoles?

A

Induced dipoles are formed in neighboring molecules due to the presence of a temporary dipole, and they are always of the opposite sign.

60
Q

What factor affects the size of London forces?

A

The number of electrons in a molecule affects the size of London forces; more electrons increase the chance of temporary dipoles forming.

61
Q

How do London forces relate to boiling points?

A

Stronger London forces require more energy to break, resulting in higher boiling points for substances with more electrons.

This explains why I2 is a solid while Cl2 is a gas.

62
Q

How does the size of molecules affect London forces in halogens?

A

The increasing boiling points of halogens down the group 7 series are due to the increasing number of electrons in larger molecules, which increases the size of London forces.

63
Q

How does the shape of a molecule influence London forces?

A

Long straight-chain alkanes have a larger surface area of contact, resulting in stronger London forces compared to spherical-shaped branched alkanes.

64
Q

What are permanent dipole-dipole forces?

A

Permanent dipole-dipole forces occur between polar molecules.

65
Q

How do permanent dipole-dipole forces compare to London forces?

A

They are stronger than London forces, resulting in higher boiling points for the compounds.

66
Q

What characterizes polar molecules?

A

Polar molecules have a permanent dipole and are asymmetrical, with significant electronegativity differences between atoms.

67
Q

Which bonds are commonly found in polar molecules?

A

Commonly found bonds include C-Cl, C-F, C-Br, H-Cl, and C=O.

68
Q

What is hydrogen bonding?

A

Hydrogen bonding occurs in compounds with a hydrogen atom attached to nitrogen, oxygen, or fluorine, which must have a lone pair of electrons.

69
Q

What is the significance of electronegativity in hydrogen bonding?

A

There is a large electronegativity difference between hydrogen and the electronegative atom (O, N, F).

70
Q

What is the bond angle in hydrogen bonding?

A

The bond angle is 180° around the hydrogen atom due to electron pair repulsion.

71
Q

How do alcohols compare to alkanes in terms of boiling points?

A

Alcohols have higher boiling points and lower volatility compared to alkanes with a similar number of electrons.

72
Q

Which substances can form hydrogen bonds?

A

Alcohols, carboxylic acids, proteins, and amides can all form hydrogen bonds.

73
Q

How many hydrogen bonds can water form per molecule?

A

Water can form two hydrogen bonds per molecule due to the two lone pairs of electrons on the oxygen atom.

74
Q

Why does ice have a lower density than liquid water?

A

In ice, molecules are held further apart by hydrogen bonds compared to liquid water.

75
Q

What is stronger than other types of intermolecular bonding?

A

Hydrogen bonding is stronger than the other two types of intermolecular bonding.

76
Q

What causes the anomalously high boiling points of H2O, NH3, and HF?

A

The anomalously high boiling points of H2O, NH3, and HF are caused by hydrogen bonding between these molecules in addition to their London forces.

77
Q

What is the effect of increasing London forces on boiling points?

A

The general increase in boiling point from H2S to H2Te or from HCl to HI is caused by increasing London forces between molecules due to an increasing number of electrons.

78
Q

What happens when an ionic lattice dissolves in water?

A

When an ionic lattice dissolves in water, it involves breaking up the bonds in the lattice and forming new bonds between the metal ions and water molecules.

79
Q

How do negative and positive ions interact with water molecules?

A

The negative ions are attracted to the δ+ hydrogens on the polar water molecules, and the positive ions are attracted to the δ- oxygen on the polar water molecules.

80
Q

Why are smaller alcohols soluble in water?

A

The smaller alcohols are soluble in water because they can form hydrogen bonds with water.

81
Q

What affects the solubility of alcohols in water?

A

The longer the hydrocarbon chain, the less soluble the alcohol.

82
Q

What is the balance in solubility of a solute in a solvent?

A

Solubility of a solute in a solvent is a complicated balance of energy required to break bonds in the solute and solvent against energy given out making new bonds between the solute and solvent.

83
Q

What types of compounds are insoluble in water?

A

Compounds that cannot form hydrogen bonds with water molecules, such as polar molecules like halogenoalkanes or non-polar substances like hexane, will be insoluble in water.

84
Q

What is the relationship between charge density and hydration enthalpy?

A

The higher the charge density, the greater the hydration enthalpy, as smaller ions or ions with larger charges attract the water molecules more strongly.

85
Q

What determines solubility in non-aqueous solvents?

A

Compounds with similar intermolecular forces to those in the solvent will generally dissolve.

86
Q

What type of solutes dissolve in non-polar solvents?

A

Non-polar solutes will dissolve in non-polar solvents, e.g., iodine dissolves in hexane.

87
Q

Why is propanone a useful solvent?

A

Propanone has both polar and non-polar characteristics, forming London forces with non-polar substances and hydrogen bonds with water.

88
Q

What is metallic bonding?

A

Metallic bonding is the electrostatic force of attraction between positive metal ions and delocalised electrons.

89
Q

What factors affect the strength of metallic bonding?

A
  1. Number of protons/strength of nuclear attraction. 2. Number of delocalised electrons per atom. 3. Size of ion.
90
Q

How does the number of protons affect metallic bonding?

A

The more protons, the stronger the bond.

91
Q

How does the number of delocalised electrons affect metallic bonding?

A

The more delocalised electrons, the stronger the bond.

92
Q

How does the size of the ion affect metallic bonding?

A

The smaller the ion, the stronger the bond.

93
Q

Why do metals have high melting points?

A

Strong electrostatic forces between positive ions and delocalised electrons require a lot of energy to break.

94
Q

Why does magnesium have a higher melting point than sodium?

A

Mg has stronger metallic bonding due to more outer shell electrons and a smaller ion size, leading to stronger electrostatic attraction.

95
Q

Why can metals conduct electricity well?

A

Delocalised electrons can move through the structure.

96
Q

Why are metals malleable?

A

Positive ions in the lattice are identical, allowing planes of ions to slide over one another easily.

97
Q

What are the four types of crystal structures?

A

Ionic, metallic, molecular, and giant covalent (macromolecular).

98
Q

What is a giant ionic lattice?

A

A structure found in ionic solids, such as sodium chloride.

99
Q

What is a giant covalent lattice?

A

A structure found in covalently bonded solids, such as diamond, graphite, and silicon(IV) oxide.

100
Q

What is a giant metallic lattice?

A

A structure found in solid metals, such as magnesium or sodium.

101
Q

What is an example of a molecular solid?

A

Iodine.

102
Q

What is the structure of diamond?

A

A giant covalent lattice with a tetrahedral arrangement of carbon atoms, each forming 4 covalent bonds.

103
Q

What is the structure of graphite?

A

A giant covalent lattice with a planar arrangement of carbon atoms, each forming 3 covalent bonds in layers.

104
Q

What is the structure of ice?

A

A molecular solid with a central water molecule in a tetrahedral arrangement, held together by ordinary covalent bonds and hydrogen bonds.

105
Q

What explains the lower density of ice compared to liquid water?

A

The molecules in ice are held further apart than in liquid water.

106
Q

What type of forces hold molecular solids together?

A

Weak London forces.

107
Q

What is the significance of strong covalent forces in giant covalent structures?

A

They result in very high melting points due to the energy required to break many strong covalent bonds.

108
Q

What is a macromolecular allotrope of carbon that cannot conduct electricity?

A

Diamond cannot conduct electricity because all 4 electrons per carbon atom are involved in covalent bonds. They are localised and cannot move.

109
Q

What is the arrangement of carbon atoms in diamond?

A

Tetrahedral arrangement of carbon atoms with 4 covalent bonds per atom.

110
Q

What are the properties of macromolecular structures like diamond and graphite?

A

Both have very high melting points due to strong covalent forces in the giant structure.

111
Q

Why can graphite conduct electricity?

A

Graphite can conduct electricity well between layers because one electron per carbon is free and delocalised, allowing electrons to move easily along layers.

112
Q

Why does graphite not conduct electricity between layers?

A

It does not conduct electricity from one layer to the next because the energy gap between layers is too large for easy electron transfer.

113
Q

What is graphene?

A

Graphene is a new substance that is one layer of graphite, with 3 covalent bonds per atom and the 4th outer electron per atom being delocalised.

114
Q

What is the structure of carbon nanotubes?

A

Carbon nanotubes have a planar arrangement of carbon atoms in layers, with 3 covalent bonds per atom in each layer.

115
Q

What allows graphene to conduct electricity well?

A

Graphene can conduct electricity well along the structure because one electron per carbon is free and delocalised, allowing electrons to move easily.

116
Q

What is the tensile strength of carbon nanotubes attributed to?

A

They have very high tensile strength because of the strong structure of many strong covalent bonds.

117
Q

How do nanotubes conduct electricity?

A

Nanotubes can conduct electricity well along the tube because one electron per carbon is free and delocalised, allowing electrons to move easily.

118
Q

What are potential uses of carbon nanotubes?

A

Nanotubes have potentially many uses, one being the potential to use as vehicles to deliver drugs to cells.

119
Q

What is present in buckminsterfullerene?

A

There are delocalised electrons in buckminsterfullerene.

Chemistry A Level (4 decks)

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