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UACE Chemistry > Chemical Bonding > Flashcards

Chemical Bonding Flashcards

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1
Q

What is a bond?

A

A bond is a force holding together two atoms or two ions in an element or compound.

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2
Q

What are the types of bonding

A
  1. Ionic bonding/Electrovalent bonding
  2. Covalent bonding
    - Normal covalent bonding
    - Coordinate bonding
  3. Metallic bonding
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3
Q

Define an Ionic bond

A

An electrostatic attraction between the electric charges of a cation (positive ion) and an anion (negative ion).

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4
Q

What does the octet rule state?

A

Elements tend to lose electrons, gain electrons or share electrons in order to acquire a noble gas core electron configuration.

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5
Q

True or false
Only valence electrons are involved in bonding except transition elements

A

True
Transition elements use both valence electrons and inner d-electrons

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6
Q

What are the physical properties of ionic compounds?

A
  • High melting and boiling point
  • Very low volatility
  • Good electrical conductivity in both molten and aqueous states bug not solid state
  • Soluble in water but insoluble in non-polar solvents
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7
Q

Why is NaCl soluble in water but not in methylbenzene?

A

There is no attraction between the ions of NaCl which is a polar component and those of methylbenzene which is non-polar thus the ions remain in the lattice.

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8
Q

What is the structure of ionic compounds?

A

Giant ionic structures/ lattices
(Regular arrangement of positively charged and negatively charged ions in regular repeating units)

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9
Q

What is the crystal structure of sodium chloride?

A

Face centered cubic

  • Each Na+ ion is surrounded by six equidistant chloride ions and each Cl- surrounded by six equidistant Na+ ions.
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10
Q

What is the crystal structure of caesium chloride?

A

Body centered cubic

  • Each Cs+ ion is surrounded by eight equidistant chloride ions.
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11
Q

Why can the caesium ion accommodate more chloride ions than the sodium ion?

A

The ionic radius of Cs+ ion is bigger than that of the Na+ ion.

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12
Q

Define coordination number of an ionic lattice

A

This is the number of ions that surround another of the opposite charge in an ionic lattice

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13
Q

Define a covalent bond

A

This is an electrostatic attraction between a shared pair of electrons and the positively charged nuclei.

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14
Q

Boron trichloride has what shape?

A

Trigonal planar

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15
Q

Ammonia has what shape?

A

Trigonal pyramidal

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16
Q

What is a co-ordinate bond (dative covalent bond) ?

A

This is a covalent bond in which both shared electrons are from the same atom.

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17
Q

What are the conditions for coordinate bonding to occur?

A

The donor atom must have a lone pair of electrons (nuleophile) and the acceptor atom must have a vacant orbital to accommodate the lone pair of electrons.

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18
Q

True or false
A coordinate bond can be shown by an arrow from the acceptor atom to the donor atom.

A

False
A coordinate bond can be shown by an arrow from the donor atom to the acceptor atom.

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19
Q

Describe the structures of covalent substances

A

Simple molecular structures
- Atoms in the molecule are joined by strong covalent bonds but their molecules are held together by weak intermolecular forces which needlittle energy to break.
- So these substances have low melting and boiling points.
- They are usually gases, liquids or waxy solids.
- They do not conduct electricity. Examples include carbon dioxide (CO2), water (H2O), sulphur S8 rings, Iodine(I2), phosphorus, P4 molecules, carbon-60 e.t.c

Giant molecular structures (covalent network solids)
- Thousands of non-metal atoms joined to each other in a regular pattern by strong covalent bonds resulting into a giant covalent lattice.
- The structure is extremely strong because of the many bonds involved. Examples include: Diamond, graphite, silicon (IV) oxide (SiO2).

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20
Q

Describe the structure of diamond

A
  • Each carbon atom is bonded to four other carbon atoms tetrahedrally by strong covalent bonds forming a giant three dimension molecular structure (macromolecular structure)
  • Each carbon atom in diamond uses all its four valence electrons for bonding.
  • Therefore diamond does not conduct electricity because it lacks free mobile electrons.
  • Diamond is very hard due to its extensive strong covalent bonds and has a high density because its atoms are closely packed.
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21
Q

Describe the structure of graphite

A
  • A crystal of graphite consists of layers of carbon atoms.
  • Each layer is a two dimension giant molecule (macromolecule) consisting of regular hexagons.
  • At the corner of each hexagon is a carbon atom joined to three other carbon atoms by strong covalent bonds.
  • Each carbon atom uses only three valence electrons for bonding.
  • Therefore graphite conducts electricity because it has free mobile electrons.
  • Graphite is soft because its layers can slide over each other due to weak intermolecular forces (Van der Waal forces).
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22
Q

Describe the structure of silicon(iv) oxide

A
  • It is a three dimension network in which each silicon atom is bonded to four oxygen atoms by single covalent bonds
  • Each oxygen atom is bonded to two silicon atoms by single covalent bonds in tetrahedral arrangement.
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23
Q

Define a polar bond

A

This is one with an unsymmetrical distribution of electron density and is represented by partial charges, δ+ and δ- .

  • The partial charges arise as a result of one atom being more electronegative than the other.
  • The more electronegative atom has a stronger attraction for the shared pair of electrons thereby attaining a partial negative charge, δ- and the less electronegative atom attaining a partial positive charge, δ+.
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24
Q

What is a non-polar bond?

A

This one with a symmetrical (an even) distribution of electron density.

  • In such bonds the two bonded atoms are of the same electronegativity therefore there are no partial charges on the bonded atoms.
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25
Q

What is a non-polar molecule?

A

A non-polar molecule is one whose net permanent dipole moment is zero.

  • Molecules whose bonds are all non-polar are definitely non-polar molecules because they have zero net permanent dipole moment.
  • Such molecules include: H2, PH3, NBr3, O2, N2 and halogens.
  • Molecules with polar bonds also become non-polar molecules if the molecule is symmetrical.
  • A molecule is symmetrical if the central atom is symmetrically surrounded by identical atoms.
  • Such molecules include: Tetrachloromethane (CCl4), dichlromethane, (CH2Cl2), boron trifluoride (BF3), carbon dioxide (CO2), methane (CH4), Benzene (C6H6), all hydrocarbon molecules.
26
Q

What is a polar molecule?

A

This is one whose net permanent dipole moment is not zero.

Examples of polar molecules include: water (H2O), ammonia (NH3), trichloromethane (CHCl3),
fluoromethane (CH3F), sulphur dioxide (SO2), nitrogen trifluoride (NF3)

27
Q

What are the three major types of intermolecular forces of attraction?

A
  1. London forces
    (also known as Dispersion forces or induced dipole- Induced dipole forces)
  2. Dipole – dipole forces
    (also known as Permanent dipole – Permanent dipole forces)
  3. Hydrogen bonding

The relative strengths of these molecular forces (other factors kept constant) are in order:
Hydrogen bonding > dipole – dipole forces > London forces

28
Q

What name is given to London forces and dipole-dipole forces together?

A

Van der Waal forces

29
Q

True or false
London forces exist in all molecules.

A

True

30
Q

How do London forces arise?

A
  • When temporary dipoles (induced dipoles) attract each other.
  • The induced dipoles result from the random distortion of the electron density distribution of one molecule due an electric field of its neighbouring molecule.
  • This causes asymmetrical distribution of the electron cloud of the molecule giving rise to temporary dipoles.
  • The formed temporary dipoles cause electron displacement in another neighbouring molecule inducing more temporary dipoles.
  • The electrostatic attraction between the opposite partial charges of the induced dipoles gives rise to London forces/ Dispersion forces.
31
Q

What factors affect the magnitude of London forces?

A

a) Number of electrons in the molecule (relative molecular mass)
- The higher the number of electrons, the further away from the nucleus the valence electrons are.
- This reduces the nuclear attraction of the valence electrons and increasing the polarizability of molecule’s electron cloud.
- The greater the polarizability, the greater is the size of the induced dipoles and the stronger the London forces.

b) Shape of the molecule
- Linear molecules (Straight chain molecules) have a larger surface area for interaction than their branched isomers.
- Straight chain isomers are therefore more closely packed giving rise to more induced dipoles and stronger London forces.
- The stronger London forces require a lot of heat break to raise enough vapour pressure equal to the atmospheric pressure for boiling to occur.
- Branched isomers have smaller areas of contact due to their more nearly spherical shape.
- They form weaker induced dipoles and weaker London forces which require less

32
Q

What is Polarizability?

A

This is the measure of the ease with which the electron cloud of an atom or molecule is distorted by an electric field

33
Q

True or false
Dipole – dipole forces exist in all polar molecules with a permanent dipole moment

A

True

These include H2O, HCl, HF, HBr and CH3COCH3.
In such molecules the positive end of the dipole of one molecule electrostatically attracts the negative end of the dipole of another adjacent molecule.

The magnitude of a dipole–dipole force depends on the size of the dipole moment of the molecule involved; the greater the dipole moment, the more polar the molecule and the stronger the dipole–dipole forces.

The boiling points of polar substances with similar molecular masses can be explained in terms of the strength of the dipole-dipole forces.

Note: Dipole–dipole forces are often called permanent dipole–dipole forces because they only occur between molecules with permanent dipole moments.

34
Q

What is a hydrogen bond?

A

This is a permanent dipole–dipole attraction between a hydrogen atom covalently bonded to a highly electronegative atom (fluorine, oxygen or nitrogen) in one molecule and another highly electronegative atom (fluorine, nitrogen or oxygen) from a different or the same molecule.

35
Q

True or false
Hydrogen bonding can only occur in compounds containing H –F, H – O and H – N bonds

A

True

These include water, ammonia, hydrogen fluoride, alcohols, primary and secondary amines, and carboxylic acids.

36
Q

What are the conditions for hydrogen bonding?

A
  • a hydrogen atom covalently bonded to oxygen, nitrogen or fluorine
  • a lone pair of electrons on the highly electronegative fluorine, oxygen or nitrogen atom.
37
Q

Why is water able to form two hydrogen bonds per molecule but hydrogen fluoride and only forms one per molecule?

A
  • Water has two hydrogen atoms and two lone pairs per molecule; so water on average forms two hydrogen bonds per water molecule.
  • Hydrogen fluoride can form only one hydrogen bond per molecule because it has only a single hydrogen atom
38
Q

Why does ammonia form a single hydrogen bond?

A

It has only a single lone pair of electrons

39
Q

What is Intermolecular hydrogen bonding?

A

This occurs between two molecules which may be same or different.

40
Q

What is Intramolecular hydrogen bonding?

A

This occurs within the same molecule.

A typical example of intramolecular hydrogen bonding occurs in 2-nitrophenol in which the hydrogen atom of the phenol group, −OH, forms intramolecular hydrogen bonds with the oxygen atom of the nitro group, −NO2.

41
Q

Explain why the melting point of 4-nitrophenol is higher than that of 2-nitrophenol

A
  • 4-nitrophenol molecules associates via intermolecular hydrogen bonding because the hydrogen atom of the phenol group,−OH, cannot form an intramolecular hydrogen bond with the oxygen atom in the nitro group,−NO2, since the two groups are so further apart.
  • The strong intermolecular hydrogen bonds need a lot of heat energy to break resulting into a higher melting point.
  • 2-nitrophenol molecules associates mainly via weaker London forces since the hydrogen atom of the phenol group,−OH and oxygen atom in the nitro group,−NO2 form an intramolecular hydrogen bond due to their close proximity.
  • The weaker Van der waal of 2-nitrophenol require less heat energy to break, hence its lower melting point.
42
Q

Why are the boiling points of ammonia, water and hydrogen fluoride anomolously higher compared to those of the hydrides of other elements in groups V, VI and VII of the Periodic Table respectively?

A
  • This because ammonia, water and hydrogen fluoride form strong intermolecular hydrogen bonds that require a lot of heat energy to break.
  • The hydrogen bonding is attributed to presence of H – N, H – O and H – F bonds in these molecules.
  • The anomolous behaviour is most marked for water. This because each water molecule is capable of forming two hydrogen bonds due to presence of two lone pairs of electrons and two hydrogen atoms per water molecule.
  • For the rest of the hydrides in groups V, VI and VI they associate via weaker Van der Waal forces (London forces + dipole-dipole forces) which increases with increase in their relative molecular masses.
  • As the Van der waal forces become stronger, the boiling points of the hydrides increase as more heat energy is required to break the increasingly stronger Van der waal forces.
43
Q

Why is water a liquid at room temperature while the rest of the hydrides of Group VI elements are gases at room tempearature?

A

Due to its high boiling point attributed to strong intermolecular hydrogen bonding, water is a liquid at room temperature while the rest of the hydrides of Group VI elements are gases at room tempearature which form weak Van der waal forces that are easily broken by room tempearture.

Due to its hydrogen bonding, the boiling points of ammonia, water and hydrogen fluoride are much higher than that of methane, CH4

44
Q

True or false
Group IV hydrides all have tetrahedral shape and are symmetrical making them non-polar molecules and having London forces as the only intermolecular forces.

A

True

45
Q

Why are the boiling points of alcohols, carboxylic acids, primary amines and secondary amines are much higher than those of alkanes of similar molecular mass?

A

Due to formation of strong intermolecular hydrogen bonds, the boiling points of alcohols, carboxylic acids, primary amines and secondary amines are much higher than those of alkanes of similar molecular mass which are linked via weak Van der waal (London forces) that require less heat energy to break.

46
Q

Why are the boiling points of primary and secondary amines higher than those of tertiary amines of comparable molecular mass?

A

Unlike primary and secondary amines, tertiary amines lack the H– N bond and therefore do not form intermolecular hydrogen bonds between their molecules.

Their molecules are linked by weaker Van der Waal forces that require less heat energy to break.

47
Q

What causes increased molecular mass of carboxylic acids when dissolved in non-polar solvents such as benzene?

A

The association of the carboxylic acid molecules via hydrogen bonding forming dimers whose molecular mass is twice the actual mass of the carboxylic acid.

48
Q

Why are many polar covalent molecules such as ammonia, methylamine, methanol, ethanol, propanol, methanoic acid, ethanoic acid, propanoic acid and propanone soluble in water?

A

Due to their ability to form strong hydrogen bonds with water molecules.

However the solubility of polar organic molecules such as amines, alcohols, carboxylic acids and carbonyl compounds in water decreases as the hydrocarbon part gets bigger since it is non-polar and water insoluble.

49
Q

Why do Carboxylic acids tend to be more water soluble than alcohols of similar molecular mass and shape?

A

Carboxylic acids form stronger hydrogen bonds with water molecules than alcohols.

The carbonyl group in carboxylic acids exerts a negative inductive effect (withdraws electrons) and this increases the size of the partial positive charge on the hydrogen atom of the −OH group allowing it form a stronger electrostaic attraction (hydrogen bond) with the oxygen atom of the water molecule wich carries a partial negative charge.

50
Q

What causes the lower density and higher volume of ice compared to liquid water?

A

Hydrogen bonding.

In ice, each water molecule tetrahedrally joined to four other water molecules through intermolecular hydrogen bonding.

The water molecules arrange themselves to maximize the number of hydrogen bonds but minimize the energy giving rise to a lattice with a large amount of space between the molecules.

The open structure of ice explains why ice has a larger volume and is less dense than water at 0°C.

When ice melts, some of the hydrogen bonds are broken allowing closer packing of the water molecules leading to a decrease in volume but an an increase in density of the water formed.

51
Q

What is a metallic bond?

A

This is the electrostatic attraction between a lattice of positive ions (cations) and delocalized electrons.

A metallic bonding is a non-directional bond.

52
Q

What are the factors affecting the strength of a metallic bond?

A
  1. The number of delocalised (valence) electrons per atom.

The greater the number of delocalised electrons per atom the stronger the electrostatic attraction between the positive metal ions and the delocalized electrons.

  1. The charge of the metal ion.

The bigger the charge on the metal ion the stronger the electrostatic attraction between the positive metal ions and the delocalized electrons.

  1. The ionic radius of the metallic action.

The smaller the ionic radius the stronger the electrostatic attraction between the positive metal ions and the delocalized electrons.

53
Q

Why is the melting point of sodium is greater than that of potassium?

A
  • Both sodium and potassium have the same number of valence electrons and charge on their ions.
  • However, the ionic radius is smaller for the sodium ion meaning that sodium has a stronger electrostatic attraction between its positive ions and the delocalized electrons than potassium.
  • Therefore a stronger metallic bond for sodium needs more heat energy to break than potassium
54
Q

Why is the melting point of calcium is greater than that of potassium?

A
  • Calcium has two delocalized electrons per atom while potassium has only one delocalized electron per atom.
  • Also the calcium ion carries a larger positive charge (2+) than the potassium ion (1+).
  • Therefore, the electrostatic attraction between the positive ions and the delocalized electrons is stronger for calcium than potassium.
  • This results into a stronger metallic bond for calcium that requires more heat energy to break than potassium.
55
Q

Describe the structure of metals

A

Metals have giant metallic structures consisting of a regular lattice arrangement of positive ions surrounded by a sea of delocalized electrons.

56
Q

True or false
All metals are good conductors of electricity due to presence delocalized electrons which carry an electric current when a potential is applied across the metal.

A

True

57
Q

True or false
The presence of impurities restricts the movement of the delocalised electrons
through the metal lattice

A

True
It results into increased electrical resistance.

This explains the copper wires used in electrical wiring need to be of a high degree of purity.

58
Q

True or false
All metals are good conductors of heat due to presence delocalized electrons
that carry heat from areas of high temperature to those of low temperature.

A

True

59
Q

Why can metals be hammered into a sheet or any other shape without breaking?

A

This is because layers of positive metal ions can slide over each other while still being held together by the delocalized eletrons (without breaking the metallic bond).

The metallic bond being non-directional means the metal can extend in any direction if hammered.

60
Q

Why do metals generally have high melting points?

A

Due stronger electrostatic attraction
between the positive metal ions and the delocalized electrons.

61
Q

What is an alloy?

A

This is typically a homogeneous mixture of metals or a mixture of a metal and a non-metal, usually carbon but sometimes phosphorus.

Alloys have improved properties compared to pure metals such as greater resistance to corrossion, greater tensile strength, greater magnetic properties and greater ductility.

62
Q

Why are alloys stronger than pure metals?

A
  • If atoms of different size or properties are introduced into the metal lattice, the regular network of positive ions is disturbed and making it more difficult for the layers of positive ions to slide over each other.
  • This is why alloys are generally, much stronger than pure metals.

UACE Chemistry (27 decks)

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