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Chem - physical > 3. Bonding > Flashcards

3. Bonding Flashcards

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1
Q

Define ionic bonding

A
  • The electrostatic force of attraction between oppositely charged ions formed by electron transfer
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2
Q

Describe how metal atoms form ions in ionic bonding

A
  • Lose electrons to form +ve ions (cations)
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3
Q

Describe how non-metal atomhs form ions in ionic bonding

A
  • Gain electrons to form -ve ions (anions)
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4
Q

Describe the structure of ionic crystals

A
  • Structure of giant lattices of ions
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5
Q

Give examples of ionic bonds

A
  • Sodium chloride
  • Magnesium oxide
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6
Q

When would ionic bonding be stronger & the melting points higher?

A
  • Ionic bonding is stronger and the melting points higher when the ions are smaller and/ or have higher charges.
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7
Q

Explain why MgO has a higher m.p. than NaCl

A
  • MgO has a higher melting point than NaCl as the ions involved (Mg2+ & O2- are smaller and have higher charges than those in NaCl , Na+ & Cl- )
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8
Q

Explain why positive ions are smaller compared to their atoms

A
  • Has one less shell of electrons
  • But same number of protons
  • The ratio of protons to electrons has increased * Greater net force on remaining electrons holding them more closely.
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9
Q

Explain why negative ions formed from groups 5-7 are larger than their atoms

A
  • The negative ion has more electrons than the corresponding atom
  • But the same number of protons.
  • So the pull of the nucleus is shared over more electrons and the attraction per electron is less, making the ion bigger.
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10
Q
A

N3- O2- F- and Na+ Mg2+ Al3+ all have the same electronic structure (of the noble gas Ne)
There are increasing numbers of protons from N to F and then Na to Al but the same number of electrons.
* The effective nuclear attraction per electron therefore increases and ions get smaller.

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11
Q

Describe what happens to ionic radii as you go down the group

A
  • Within a group the size of the ionic radii increases going down the group.
  • This is because as one goes down the group the ions have more shells of electrons.
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12
Q

Define a covalent bond

A
  • Shared pair of electrons attracted to the nuclei of both bonding atoms
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13
Q

State the two possible structures of covalent bonds

A
  • simple molecular (simple covalent structure)
  • macromolecular (giant covalent structure)
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14
Q

Give examples of simple covalent structures

A
  • Iodine
  • Ice
  • Carbon dioxide
  • Water
  • Methane
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15
Q

Give examples of macromolecular structures

A
  • Diamond
  • Graphite
  • Silicon dioxide
  • Silicon
  • Graphene
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16
Q

Describe the properties of a simple molecular atom

A
  • Mostly gases and liquids
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17
Q

Describe the boiling and melting point of simple molecular atoms

A
  • Low b.p. and m.p.
  • Because of weak intermolecular forces between molecules
  • (specify type e.g van der waals/hydrogen bond)
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18
Q

Describe the solubility of simple molecular atoms

A
  • Generally poor
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19
Q

Describe the conductivity of simple atoms

A
  • Poor when solid & molten
  • No ions to conduct and electrons are localised (fixed in place)
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20
Q

Describe the properties of a macromolecular atom

A
  • Solids
  • Giant molecular structure
    With intermolecular forces (van der Waals, permanent dipoles, hydrogen bonds) between molecules??
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21
Q

Describe the boiling and melting point of macromolecular atoms

A
  • High melting and boiling point
  • Because of many strong covalent bonds in macromolecular structure.
  • Take a lot of energy to break the many strong bonds
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22
Q

Describe the solubility macromolecular atoms

A
  • Insoluble
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23
Q

Describe the conductivity macromolecular atoms

A
  • Diamond and sand: poor, because electrons can’t move (localised)
  • Graphite: good as free delocalised electrons between layers
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24
Q

Define dative covalent bonding

A
  • Where both electrons come from one atom, a pair of electrons has been donated
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25
Q

When does a dative covalent bond form?

A
  • When the shared pair of electrons in the covalent bond come from only one of the bonding atoms.
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26
Q

What is dative covalent bonding also known as?

A
  • Co-ordinate bonding.
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27
Q

Give common examples of dative covalent bonds

A
  • NH4+
  • H3O+
  • NH3BF3
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28
Q

How would you prove that a bond is a dative covalent bond e.g. between BF3?

A
  • Lone pair of electrons donated from fluoride to BF3
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29
Q
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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30
Q
A

The dative covalent bond acts like
an ordinary covalent bond when
thinking about shape so in NH4+ the shape is tetrahedral

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31
Q

Define metallic bonding

A
  • The electrostatic force of attraction between the positive metal ions and the delocalised electrons
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32
Q

Name the three main factors that affect the strength of metallic bonding

A
  • Number of protons/ Strength of nuclear attraction.
    The more protons the stronger the bond
    1. Number of delocalised electrons per atom (the outer shell electrons are delocalised)
      The more delocalised electrons the stronger the bond
    1. Size of ion.
      The smaller the ion, the stronger the bond.
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33
Q

Explain why Mg has stronger metallic bonding than Na

A

hence a higher melting point. The metallic bonding gets stronger because in Mg there are more electrons in the outer shell that are released to the sea of electrons. The Mg ion is also smaller and has one more proton. There is therefore a stronger electrostatic attraction between the positive metal ions and the delocalised electrons and higher energy is needed to break bonds.

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34
Q

Describe the structure of metallic bonding

A
  • Giant metallic lattice
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35
Q

Describe the boiling and melting point of metallic bonds

A
  • High - strong electrostatic forces between positive ions and sea of delocalised electrons
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36
Q

Describe the solubility of metallic bonds

A
  • Insoluble
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37
Q

Describe the conductivity of metallic bonds

A
  • Good - delocalised electrons can move through structure
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38
Q

Describe the properties of metallic bonds

A
  • Shiny metal
  • Malleable as the positive ions in the lattice are all identical.
  • So the planes of ions can slide easily over one another
    -attractive forces in the lattice are the same whichever ions are adjacent
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39
Q

Define electronegativity

A
  • The relative tendency of an atom in a covalent bond in a molecule to attract electrons in a covalent bond to itself.
    OR
  • The power of an atom to attract the electron density of a covalent bond to itself
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40
Q

Name the most electronegative atoms

A
  • F, O, N, and Cl
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41
Q

Name the most electronegative element

A
  • Fluorine
  • Value of 4.0
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42
Q

How is electronegativity measure?

A
  • On the Pauling scale (ranges from 0-4)
43
Q

Name the factors affecting electronegativity

A
  • Number of protons
  • Distance
  • Shielding
44
Q

Describe how the number of protons affects the electronegative of an atom

A
  • As no. protons increase, electronegativity increases
45
Q

Explain why electronegativity increases across a period

A
  • Electronegativity increases across a period as the number of protons increases
  • And the atomic radius decreases
  • Because the electrons in the same shell are pulled in more.
46
Q

Explain why electronegativity decreases down a group

A
  • It decreases down a group because the distance between the nucleus and the outer electrons increases and the shielding of inner shell electrons increases
47
Q

What are the extremes of a continuum of bonding type?

A
  • Ionic and covalent bonding are the extremes of a continuum of bonding type.
48
Q

What can the difference electronegative between elements determine?

A
  • Where a compound lies on the scale of ionic and covalent bonding
49
Q

How would a compound be purely covalent?

A
  • If a compound contained elements of similar electronegativity and hence a small electronegativity difference
50
Q

How would a compound be ionic?

A
  • If a compound contained elements of very different electronegativity and hence a very large electronegativity difference (> 1.7)
51
Q

Define polarity

A
  • Unequal sharing of electrons in a bond
52
Q

When would a polar covalent bond form? (permanent dipole)

A
  • When the elements in the bond have different
    electronegativities . (Of around 0.3 to 1.7)
53
Q

Describe what happens when a bond is a polar covalent bond

A
  • It has an unequal distribution of electrons in the bond and produces a charge separation, (dipole)
  • δ+ δ- ends.
54
Q

How are two atoms that are bonded together with different electronegativity represented ?

A
  • Using partial charges
  • Delta positive - electron deficient
  • Delta negative - electron rich
55
Q

Which element will be the δ- end in a polar compound?

A
  • The element with the larger electronegativity
56
Q

Give an example of a polar molecule

A
  • δ+ δ–
    H – Cl
57
Q

Describe a symmetrical molecule in terms of its polarity

A
  • A symmetric molecule (all bonds identical and no lone pairs) will not be polar even if individual bonds within the molecular are polar.
58
Q

Give an example of a symmetrical molecule (non-polar)

A
  • CO2
  • O = C = O
59
Q

When would a molecule be non polar?

A
  • When electron sharing is equal due to the symmetrical shape of the molecule
  • The individual dipoles on the bonds ‘cancel out’
  • There is no net dipole moment
60
Q

Compare the polarity of CCl4 & CH3Cl

A
  • CCl4 will be non-polar whereas CH3Cl will be polar
61
Q

Define intramolecular forces

A
  • Forces within a molecule (usually covalent bonds)
62
Q

Define intermolecular forces

A
  • Forces between molecules
63
Q

Define intermolecular forces

A
  • Forces between molecules
64
Q

Where do Van Der Waals’ forces occur?

A
  • Van der Waals forces occur between all molecular substances (simple covalent molecules) and the separate atoms noble gases.
  • They do not occur in ionic substances.
65
Q

What are Van Der Waal’s forces also known as?

A
  • These are also called transient, induced dipole-dipole interactions.
66
Q

Explain the induced dipole-dipole interactions (Van Der Waals’ forces)

A
  • In any molecule the electrons are moving constantly and randomly producing a changing dipole in a molecule
  • As this happens the electron density can fluctuate and parts of the molecule become more or less negative (i.e. small temporary or transient dipoles form.)
  • At any time an instantaneous dipole will exist but its position constantly changes (temporary dipole)
  • These instantaneous dipoles can cause dipoles to form in neighbouring molecules.
  • Electrons continually move so are created and destroyed all the time, overall atoms are attracted to one another
  • The induced dipole is always the opposite sign to the original one.
67
Q

What are the main factor affecting size of Van Der Waals’ forces?

A
  • The more electrons there are in the molecule the higher the chance that temporary dipoles will form.
  • This makes the Van der Waals stronger between the molecules and so boiling points will be greater
68
Q

How can the increasing b.p. of the halogens down group 7 be explained using VDW forces?

A
  • By the increasing number of electrons in the bigger molecules causing an increase in the size of the Van der Waals between the molecules.
  • This is why I2 is a solid whereas Cl2 isagas.
69
Q

How can the increasing b.p of the alkane homologous series be explained using VDW forces?

A
  • Increasing number of electrons in the bigger molecules causing an increase in the size of the Van der Waals between molecules.
70
Q

How does the shape of the molecule effect the size of the Van Der Waals’ forces?

A
  • Long chain alkanes have a larger surface area of contact between molecules for Van der Waals to form than compared to spherical shaped branched alkanes
  • So have stronger Van der Waals.
71
Q

Where do permanent dipole-dipole forces occur?

A
  • Permanent dipole-dipole forces occurs between polar molecules
72
Q

Which is stronger, Van der Waals or permanent dipole-dipole forces?

A
  • Permanent dipole-dipole so the compounds have higher boiling points
73
Q

Describe the characteristics of polar molecules

A
  • Polar molecules have a permanent dipole. (commonly compounds with C-Cl, C-F, C-Br H-Cl, C=O bonds)
  • Polar molecules are asymmetrical and have a bond where there is a significant difference in electronegativity between the atoms.
74
Q

Do permanent dipole-dipole forces occur in addition to Van der Waals forces?

A
  • True
75
Q

Define hydrogen bonding

A
  • A special type of permeant dipole - permenant dipole
75
Q

State the compounds which can form hydrogen bonds

A
  • Alcohols, carboxylic acids, proteins, amides
76
Q

When does hydrogen bonding occur?

A
  • In compounds that have a hydrogen atom attached to one of the three most electronegative atoms of nitrogen, oxygen and fluorine, which must have an available lone pair of electrons.
77
Q

Give an example of hydrogen bonding

A

e.g. –O-H, -N-H, F- H bonds

78
Q

Why does hydrogen bonding occur between H and O,N, F?

A
  • There is a large electronegativity difference between the H and the O,N,F
79
Q

What must you always do when drawing out hydrogen bonds?

A
  • Always show the lone pair of electrons on the O,F,N
  • And the dipoles and all the δ- δ+ charges
80
Q

Do hydrogen bonding occur in addition to Van der Waals’ forces?

A
  • True
81
Q

Which is the strongest type of intermolecular bonding?

A
  • Hydrogen bonding
82
Q

Why does H2O, NH3 and HF have an anomalously high b.p.?

A
  • Caused by the hydrogen bonding between the molecules
83
Q

What is the general increase in b.p. from H2S and H2Te caused by?

A
  • Increasing Van der Waals forces between molecules due to an increasing number of electrons.
84
Q

State the four type of crystal structure

A
  • ionic, metallic, molecular and giant covalent (macromolecular).
85
Q

Describe the structure of ionic crystals

A
  • Giant Ionic lattice showing alternate positive and negative ions
  • Cube shaped
86
Q

Describe the structure of iodine (molecular crystal)

A
  • Regular arrangement of I2
  • 2 molecules held together by weak van der Waals forces
87
Q

Describe the structure of diamond (macromolecular crystal)

A
  • Tetrahedral arrangement of carbon atoms
  • 4 covalent bonds per atom
88
Q

Describe the structure of graphite (macromolecular crystal)

A
  • Planar arrangement of carbon atoms in layers. 3 covalent bonds per atom in each layer.
  • 4th outer electron per atom is delocalised. * Delocalised electrons between layers.
89
Q

Why do macromolecular structures have very high m.p?

A
  • Because of strong covalent forces in the giant structure.
  • It takes a lot of energy to break the many strong covalent bonds.
90
Q

E.g. In terms of structure and bonding, explain why the b.p. of bromine is different from magnesium

A
  • STRUCTURE: Bromine is simple molecular, whereas magnesium is metallic (positive ions in a sea of delocalised electrons
  • STRENGTH: Br2 has weak van der Waals forces between molecules, Mg has strong metallic bonds, more energy needed to overcome
  • Mg has greater liquid range bc forces of attraction in liquid metal are stronger
91
Q

E.g. Explain why m.p of iodine is low, and hydrogen iodide very low

A
  • TYPE: I2, and HI is molecular
  • FORCES: IMF hold molecules together, they are weak hence low m.p.
  • SIZE: I2 bigger than HI, so has more electrons
  • Therefore stronger VDW between molecules (more electrons stronger VDW) so it requires more energy to break, m.p. higher
  • NO EFFECT: HI shows permanent dipole-dipole attraction, but less than VDW forces in iodine
92
Q

Describe the shape of linear molecules

A
  • No. bonding pairs : 2
  • No. lone pairs: 0
  • Bond angle: 180
  • Examples:
93
Q

Describe the shape of trigonal planar molecules

A
  • No. bonding pairs : 3
  • No. lone pairs: 0
  • Bond angle: 120
  • Examples:
94
Q

Describe the shape of tetrahedral molecules

A
  • No. bonding pairs : 4
  • No. lone pairs: 0
  • Bond angle: 109.5
  • Examples:
95
Q

Describe the shape of trigonal pyramidal molecules

A
  • No. bonding pairs : 3
  • No. lone pairs: 1
  • Bond angle: 107
  • Examples:
96
Q

Describe the shape of bent molecules (non-linear)

A
  • No. bonding pairs : 2
  • No. lone pairs: 2
  • Bond angle: 104.5
  • Examples:
97
Q

Describe the shape of trigonal bipyramidals

A
  • No. bonding pairs : 5
  • No. lone pairs: 0
  • Bond angle: 120 and 90
  • Examples:
98
Q

Describe the shape of octahedral molecules

A
  • No. bonding pairs : 6
  • No. lone pairs: 0
  • Bond angle: 90
  • Examples:
99
Q

How do you differ between trigonal pyramidal & trigonal planar? (linear & bent)

A
  • Look at the main element of the compound, see how many electrons it has
  • If all of its electrons is being bonded e.g. BF3 then no lone pairs, however if it has electrons that still need bonding, it has a lone pair.
100
Q

Explain how you would answer a question regarding molecule shape

A
  1. State number of bonding pairs and lone pairs of electrons.
  2. State that electron pairs repel and try to get as far apart as possible (or to a
    position of minimum repulsion.)
  3. If there are no lone pairs state that the electron pairs repel equally
  4. If there are lone pairs of electrons, then state that lone pairs repel more than
    bonding pairs.
  5. State actual shape and bond angle.
101
Q

How do lone pairs differ from bonding pairs?

A
  • Lone pairs repel more than bonding pairs and * Reduce bond angles (by about 2.5o per lone pair in above examples)
102
Q

Occasionally more complex shapes are seen that are variations of octahedral and trigonal bipyramidal where some of the bonds are replaced with lone pairs. You do not need to learn the names of these but ought to be able to work out these shapes using the method below

A
103
Q

add more square planar etc….

A

Chem - physical (12 decks)

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