Brainscape's Knowledge GenomeTM


Top Flashcards (Certified)
All Flashcards

AS Chemistry > 2 - Bonding & Structure > Flashcards

2 - Bonding & Structure Flashcards

(72 cards)

Start Studying
1
Q

How are ions held together and why?

A

Giant Ionic Lattices, maximising attractive force and minimising repulsive. The lattice itself is made up of strong electrostatic forces between ions of opposite charge.

How well did you know this?
1
Not at all
2
3
4
5
Perfectly
2
Q

What 3 things affect the strength of ionic bonds and the melting/boiling points of ions?

A
  • Ionic charges: greater charge = stronger ionic bond = higher MP/BP
  • Ionic radius: smaller = closer = stronger electrostatic forces = higher MP/BP
  • Geometry: if ions are touching more ions of the opposite charge, this increases stability and electrostatic forces, increasing MP/BP (e.g. a 6:6 coordinated lattice such as NaCl has lower MPs than 8:8 coordinated CsCl)
How well did you know this?
1
Not at all
2
3
4
5
Perfectly
3
Q

What are isoelectronic ions?

A

Different atoms with the same electronic configuration or number of electrons.

How well did you know this?
1
Not at all
2
3
4
5
Perfectly
4
Q

Why is the ionic radius in non-metals larger than the atomic radius?

A

Because the effective attraction between the positive nucleus and the negative electrons is weakened, allowing electrons to move further away from nucleus and providing larger ionic radius.

How well did you know this?
1
Not at all
2
3
4
5
Perfectly
5
Q

What are the trends of the ionic radius?

A
  • Increases down groups due to more energy levels
  • Decreases across periods as electrons experience greater nuclear charge effect, pulling them closer to the nucleus and decreasing the radius.
How well did you know this?
1
Not at all
2
3
4
5
Perfectly
6
Q

What are 3 examples that act as evidence for Ionic Bonding?

A
  • High MPs - ions are strongly attracted, requiring more energy to break bonds (showing positive and negative ion attraction)
  • Solubility - soluble in water but not in non-polar solvents (showing ions are charged)
  • Conductivity - not conductive when solid but do when molten (showing ions are fixed in a strong lattice)
  • Unshapeable - brittle, allowing no ion overlap due to strong repulsive forces which would break lattice (shows positive/negative charges)
  • Electron Density Maps - ionic compounds such as NaCl may have 0 electron density between ions, showing them as individual atoms without electron sharing (supports idea of lattice)
How well did you know this?
1
Not at all
2
3
4
5
Perfectly
7
Q

What is an electron density map?

A

X-Rays pass through a crystal which scatters the radiation to obtain a diffraction pattern. The amound of scattering depends on electron density, allowing you to view the position of atoms or ions within a solid.

How well did you know this?
1
Not at all
2
3
4
5
Perfectly
8
Q

What is covalent bonding?

A

The strong electrostatic attraction between the nuclei and shared electrons of 2 atoms, usually between non-metals.

How well did you know this?
1
Not at all
2
3
4
5
Perfectly
9
Q

“Atoms share pairs of electrons to reach the electronic configuration of a Noble Gas.” In what ways is this assertion correct or incorrect?

A

CORRECT:
- Halogens are 1e away from NG config; form 1 covalent bond
- Oxygen is 2e short; forms 2 covalent bonds
- Nitrogen is 3e short; forms 3 covalent bonds
- Carbon is 4e short; forms 4 covalent bonds
INCORRECT:
- Works for P2 except for Boron and for P3/4 but NOT PCl5, SF6 and Noble Gas compounds such as XeF4.

How well did you know this?
1
Not at all
2
3
4
5
Perfectly
10
Q

What is a more accurate theory to explain covalent bonds?

A

Covalent bonds are caused by atomic orbital overlap.

How well did you know this?
1
Not at all
2
3
4
5
Perfectly
11
Q

Why can Phosphorus form 5 covalent bonds? (PCl5)

A

After forming PCl3, P has 5 empty 3d orbitals, similar in energy to 3s orbitals. The 3s electrons are therefore promotd to empty 3d orbitals, giving P 5 unpaired electrons which therefore allow it to form 5 covalent bonds in an excited state.

How well did you know this?
1
Not at all
2
3
4
5
Perfectly
12
Q

What 2 factors determine the strength of covalent bonds?

A
  • Atomic radius: smaller = stronger bonds

- Number of electrons shared: more = stronger bond

How well did you know this?
1
Not at all
2
3
4
5
Perfectly
13
Q

What effect does higher electron density have on bond enthalpy and length?

A

Causes stronger attraction between atoms, creating higher bond enthalpy that in turn shortens the bond length to pull the nuclei closer together in a stronger bond.

How well did you know this?
1
Not at all
2
3
4
5
Perfectly
14
Q

What is dative covalent bonding?

A

When both electrons in a bond come from the same atom.

How well did you know this?
1
Not at all
2
3
4
5
Perfectly
15
Q

What is an example of dative covalent bonding?

A
  • NH4+
  • PCl6-
  • Al2Cl6: 2AlCl3 molecules join together. The lone pair on Cl bonds with empty Al orbital to give a dative bond. Energy is released as the dative bond is formed, making Al2Cl6 more stable than AlCl3.
How well did you know this?
1
Not at all
2
3
4
5
Perfectly
16
Q

What are the shapes of molecules and ions determined by?

A

Valence Shell Electron Pair Repulsion Theory (VSERPR): electron pairs rpel each other, so they are separated as far as possible with minimum repulsion.

How well did you know this?
1
Not at all
2
3
4
5
Perfectly
17
Q

What is the bond angle and name for a compound with 2 bonding pairs?

A

Linear: 180 bond angle

Cl - Be - Cl

How well did you know this?
1
Not at all
2
3
4
5
Perfectly
18
Q

What is the bond angle and amount of bond pairs for Trigonal Planar?

A 
3 bond pairs, 120 angle
Cl         Cl
   \  Al  /
       |
      Cl
How well did you know this?
1
Not at all
2
3
4
5
Perfectly
19
Q

How many bonding pairs does a molecule with a 109.5 bond angle have, and what is the name of this shape?

A

4 bonding pairs: Tetrahedral

How well did you know this?
1
Not at all
2
3
4
5
Perfectly
20
Q

What shape is a molecule with 5 bonding pairs, and what is the bond angle for this?

A

Trigonal Bipyramidal, 120 and 90 bond angles

How well did you know this?
1
Not at all
2
3
4
5
Perfectly
21
Q

How many bonding pairs are there in an octahedral molecule and what is the bond angle?

A

6 bonding pairs, 90 angle

How well did you know this?
1
Not at all
2
3
4
5
Perfectly
22
Q

What shape is a molecule with 3 bonding pairs and 1 lone pair, and what is its bond angle?

A

Pyramidal, 107 angle

How well did you know this?
1
Not at all
2
3
4
5
Perfectly
23
Q

How many bonding and lone pairs does a molecule with a bond angle of 104.5 have and what is this shape called?

A

2 bonding pairs, 2 lone pairs. Angular

How well did you know this?
1
Not at all
2
3
4
5
Perfectly
24
Q

What is the one molecule that is Square Planar, how many bond/lone pairs are there, and what is the bond angle?

A

XeF4

4 bond pairs, 2 lone pairs. 90 bond angle

How well did you know this?
1
Not at all
2
3
4
5
Perfectly
25
What is the shape of carbon dioxide, a molecule with 2 double bonds?
Linear - 2 bond pairs, 0 lone pairs, 180 angle | O = C = O
26
How can you predict the shapes of molecules?
1. ) Work out how many electrons are around the central atom: +1 to the group number for every bond formed. If it's an ion, add or remove electrons depending on the charge. e. g. NH3: N is central atom in G5, forms 3 bonds (5+3=8) so 8 electrons 2. ) Draw it out and count the number of lone pairs - i.e. any electrons not in a bond e. g. NH3 has 1 lone pair (3 bonds = 6 electrons, 8 - 6 = 2 which is 1 pair) 3. ) Use the amount of bond pairs and lone pairs to determine the shape (hint: memorise the shape info first) e. g. 3 bond pairs, 1 lone pair: NH3 is pyramidal (bond angle = 107)
27
What type of electrons have the biggest repulsion?
Lone + Lone | then Lone + Bond, Bond + Bond
28
How do more lone pairs affect the bond angle?
More lone pairs = smaller angle
29
What are 3 of the allotropes of carbon?
- Diamond - Graphite - Graphene - Fullerene - Carbon Nanotubes
30
What are the properties of graphite?
- Each carbon bonded to 3 carbons by sigma bonds in a layered structure (forming rings of 6C) - High MP - strong C-C bonds - Conductive - 4th e in p orbital forms delocalised electron clouds, able to carry charge through sheets - Rings held together by London forces - weak, allowing sheets to slide over one another - Insoluble - no atraction between solutes and carbons - Lower density than diamond - spaces betwen layers
31
What are the properties of graphene (C60)?
- 1 atom thick layer of graphite - Conducts both heat and electricity - Each carbon has 3 sigma bonds (delocalised 4th e) - Shape: buckyball
32
What are the properties of diamond?
- Each carbon attached to 4 more carbons by sigma bonds in a tetrahedral arrangement - High MP - strong C-C covalent bonds - Unconductive - no delocalised electrons - Insoluble - no attractions between solutes and carbon atoms
33
What are the properties of fullerene (C70)?
- Chemically similar to alkenes - 25 hexagons - Shaped like rugby ball
34
What are the properties of carbon nanotubes and what are they used for?
- Cylinders with inteconnecting carbon hexagons, sealed within fullerene - Conductive - delocalised electrons - Clusters 100x stronger than steel - Used to deliver drugs to certain body parts
35
How can the allotropes of carbon be used in medicine?
- Carbon nanotubes deliver drugs to certain body parts - Radon-224 sealed in graphene and coated in tumour-targeting antibodies for cancer treatment. Radon emits alpha rays to destroy cancer withiout causing excessive bodily harm
36
What are 4 ways you can tell if covalent bonding is present?
- High MPs (strong covalent bonds) - Hard ( strong bonds/lattice arrangement) - Insoluble (atoms not attracted to solutes) - Unconductive, except for graphite and its variations (no delocalised electrons)
37
How are metallic bonds formed?
Metal atoms lose their outer electrons to form positive cations arranged in a lattice, surrouded by sea of delocalised electrons (enabling conduction)
38
What does the strength of metallic bonding depend upon?
- Charge on metal ion (the same as the amount of delocalised electrons) - Metallic radius (larger = weaker bond) - Structure of metallic lattice
39
What are three physical properties of compounds with metallic bonding?
- High MPs: strong bonds require more energy to break - Conductive of electricity and heat: sea of delocalised electrons pass kinetic energy to one another to create heat - Malleable (shapeable) and Ductile (can be pulled into wire): layers slide over each other due to delocalised electrons preventing repulsion between layers - Insoluble (except in liquid metal): strong bonding
40
What are the trends in melting point across the periodic table?
- Increases across periods as metallic radius decreases | - Decreases down groups as radius increases (little attraction between ions and outer electrons)
41
What are the trends in group 1 and 2 metals as a result of metallic bonding?
G1: low density, low MP (only 1 electron in bond), 8 coordinated (not tightly packed or forming as many bonds), large radius (weakens bond) G2: higher density, higher MP (2 electrons in bond), smaller radius (strengthens bond), higher charge
42
A substance has a melting point of 1045K and is an insulator when solid. When melted, it conducts electricity. How is this substance bonded?
> High melting point: not simple molecular > Unconductive when solid: not metallic > Conductive when molten: not giant covalent Therefore: IONIC
43
What is electronegativity?
The ability of an atom to attract the shared electron pair in a covalent bond, measurable using the Pauling scale.
44
What is the trend of electronegativity across the periodic table and why?
- Increases across periods - Decreases down groups (except Noble Gases) > Higher electronegativity leads to increased nuclear charge, reducing the atomic radius of the atom.
45
When and how are polar bonds formed?
Between two covalently bonded elements with a large difference in electronegativity. Both atoms have a partial opposite charge, creating 2 poles (positive and negative).
46
Which region of a polar bond has the highest electron density and why?
The negative pole, because it has a greater affinity for electrons. It becomes negative as electrons congregate there.
47
Why are homonuclear diatomic gases such as Oxygen and Hydrogen non-polar?
Both atoms are equally electronegative, the charges cancelling each other out and meaning there is no polarity. However, these can still form polar bonds with other elements.
48
What types of bond would the following examples have: - two elements with same electronegativity? - two elements with slightly different electronegativity? - two elements with significantly different electronegativity?
- pure, non-polar covalent bond - polar covalent bond - ionic bond
49
How can you predict what type of bonding will occur between 2 atoms?
1. ) Look up electronegativity values on Pauling scale 2. ) Find the difference in electronegativity by subtracting the smaller value from the larger 3. ) Look up this value on the given table to find the % ionic character.
50
What must the electronegativity difference be if the bond is to be considered as polar?
Over 0.4.
51
Which of these shapes are polar and why/why not: - Tetrahedral? - Linear? - Angular? - Symmetrical?
- Tetrahedral - polar only if one of the adjoining atoms is different/replaced by a lone pair - Linear - not polar as the polarities cancel - Angular - polar, due to bond angle creating poles - Symmetrical molecules - not polar, even if the bonds are, because polarities cancel.
52
What is a simple test for the presence of polar molecules?
Rub a rod with a duster and hold it next to a stream of water from a tap. A polar substance will go TOWARDS the rod and a non-polar will do nothing.
53
How is polarity measured?
Using a molecule's dipole moment (charge difference x distance between atoms). If the dipole moments cancel, the molecule isn't polar.
54
How are London forces created?
Electrons in covalent molecules fluctuate within bonds or orbitals, motion causing a small temporary dipole in the molecule (i.e. electrons more likely to be at one pole than another at any given point). This instantaneous dipole can then induce a dipole on a nearby atom, the electrons attracting to the original atom. Because the first atom's electrons are fluctuating, the second atom's do too, in turn inducing dipole in another atom. Pretty soon, all these atoms are held together by these London forces (aka instantaneous dipole-induced dipole bonds)
55
What does the strength of London forces depend upon?
- The number of electrons in the molecule - more electrons = larger electron clouds - Surface area - greater = stronger, as more of the electron cloud is exposed - Shape - long thin molecules develop bigger dipoles and can be closer together, strengthening London forces
56
Why do the boiling points of the Noble Gases increase down the group?
Increasing number of electrons results in a larger atomic radius, this strengthening the London forces. Stronger forces require more thermal energy to overcome, thus increasing both melting and boiling points.
57
Why does pentane have a higher MP/BP than ethane?
Pentane has a longer carbon chain, resulting in stronger London forces as there is more molecular surface contact and more electrons to form the instantaneous dipoles with. The stronger forces require more thermal energy to break, hence the MP and BP is higher.
58
Why does 3,5-diethylhexane have a lower boiling point than octane?
Because it contains branched chains, meaning there is small molecular surface contact as the molecules cannot fit closely together. This in turn limits the amount of London forces that can form, reducing the boiling point of the compound.
59
What are permanent dipole-dipole interactions?
The slightly +/- charges on polar molecules cause weak electrostatic forces of attraction between molecules. These occur as well as London forces, though are significantly weaker. Molecules with permanent dipole-dipole interactions will have higher MP/BPs as there are more forces to break.
60
Which elements form hydrogen bonds and why?
Fluroine, Oxygen and Nitrogen. This is because they are very electronegative , attracting the electrons away from hydrogen to leave H+ (and make the bond polar).
61
What are the relative strengths of the hydrogen bonds formed with F and O?
- With O: stronger than with F as more H bonds per molecule in water. BP = 100 - With F: weaker than with O as only one lone pair used in bond. BP = 20
62
What is the structure of ice and why is it so strong?
A simple molecular structure containing interlocking rings of 6 H20 molecules, held together by H bonds - the arrangement allows for the maximum number of H bonds to be formed. Less dense than liquid due to the large gaps in the ring structure, allowing ice to float. As it melts, density increases as the ring structure breaks.
63
Why does Chlorine not engage in hydrogen bonding?
Despite having a high enough electronegativity to do so, its radius is too large to get close enough. As a result, it can only form dipole-dipole interactions.
64
What effect do hydrogen bonds have on melting and boiling points?
Increases them - bonds stronger than both London and dipole-dipole, requiring more energy to break.
65
What must occur in order for solutess to dissolve?
- Bonds in solute must break - Bonds in solvent must break, solvent molecules surrounding solute - New bonds must form between solute and solvent - These new bonds must have a strength EQUAL TO or GREATER THAN the original broken bonds.
66
What is a saturated solution?
A solution containing as much solute as possible at a certain temperature.
67
What is solubility?
A measure of the concentration of the saturated solution of solute at a certain temperature (moles per 100g water at 25 degrees)
68
What are the two types of solvent and what are their properties?
- Polar: made of polar molecules, most of them able to form H bonds (and so any compound able to form hydrogen bonds will dissolve in it). - Non-polar: depends on the similarities between the intermolecular forces of solute and solvent whether or not it will dissolve.
69
Why are alkanes insoluble in water but soluble in other alkanes?
Alkanes contain London forces which are identical strengths in both solvent and solute, allowing for the bonds to be broken and the new bonds formed to equal the strength of those previously broken.
70
Why are most ionic compounds soluble in water?
The ions are attracted to the oppositely charged ends of water molecules, pulling the ions away from one another in a process known as hydration. Hydration energy is released as the hydrated complexes form.
71
Are alcohols soluble in water and why/why not?
Yes - they have polar -OH group which is attraced to the O-H bond in water. However, solubility decreases with more carbons in the chains as the carbons are not soluble.
72
Why are halogenoalkanes, containing polar bonds, insoluble in water?
Despite their polar bonds, their are insoluble in water due to their weak dipoles being unable to form hydrogen bonds with water molecules. Water's hydrogen bonding to itself is much stronger than any bonds formed with the halogenoalkane, this preventing the solute from dissolving. However, halogenoalkanes can dissolve in polar substances that form the same permanent dipole-dipole bonds as they do.

AS Chemistry (5 decks)

Brainscape helps you reach your goals faster, through stronger study habits.
© 2025 Bold Learning Solutions. Terms and Conditions