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the underlying truth of chemistry explored by me > Topic 2 - Bonding > Flashcards

Topic 2 - Bonding Flashcards

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1
Q

What are ions?

A

Charged particles that is formed when an atom loses of gains electrons

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2
Q

What are molecular ions?

A

Covalently bonded atoms that lose or gain electrons

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3
Q

What is ionic bonding?

A

Strong electrostatic attraction between oppositely charged ions.

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4
Q

What is the arrangement of the ions in an ionic compound ?

A

In a giant ionic lattice.

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5
Q

What determines the strength of an ionic bond?

Name 3 things

A
  1. Ionic radius
  2. Ionic charge
  3. The number of shells in a certain element.
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6
Q

Is the electrostatic interaction in an Ionic compound directional?

A

No: all that matters is the distance between the 2 ions, not their orientation with respect to one another.

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7
Q

Explain the trend in ionic radius down a group?

A

Ionic radii increases going down the group. This is because down the group the ions have more shells if electrons and thus outermost electrons experiences less pull from positive nucleus.

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8
Q

What properties do metals have

A 
High melting temperatures 
Good electrical conductivity 
Good thermal conductivity 
Malleability 
Ductile
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9
Q

How do metals carry electric current

A

Metal have delocalised electrons which are free and mobile. When potential difference is applied across the ends of a metal, the delocalised electrons will be attracted to and move towards the positive terminal of the cells. This movement of electric charge constitutes an electric current.

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10
Q

What determines he melting temperature of a metal?

A

The number of delocalised electrons per cation.

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11
Q

Why do metals in the d block typically have higher melting temperatures?

A

Because they have more delocalised electrons per ion.

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12
Q

What two factors contribute to the ability of metals to transfer heat energy?

A
  1. The free- moving delocalised electrons pass kinetic energy along the metal.
  2. The cations are closely packed and pass kinetic energy from one cation to another.
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13
Q

Why are metals malleable and ductile?

A

Both of these properties are dependent on the ability of the delocalised electrons and the cations to move throughout the structure.
When a stress is applied to a metal, the layers of cations may slide over each other.
However since the delocalised electrons are free moving, they move with the cations and prevent repulsion between the cations in one layer and in another layer.

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14
Q

Do solid ionic compounds conduct electricity?

A

No because there are no delocalised electrons and the ions are not free to move.

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15
Q

How do molten ionic compounds conduct electricity?

A

Since the ions are mobile they will migrate to the electrodes of the opposite sign when potential difference is applied. If direct current is used, the compound will undergo electrolysis as the ions are discharged at the electrodes.

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16
Q

Why are ionic compounds soluble in water?

A

Ionic compounds are soluble in water because energy needed to break apart the lattice structure and separate the ions can be supplied by the hydration of the separated ions produced. Both positive and negative ions are attracted to water molecules because of the polarity that water molecules posses.

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17
Q

Name the physical properties of ionic compounds?

A

High melting temperatures
Brittleness
Poor electrical conductivity when solid but good when molten
Often soluble in water

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18
Q

Explain the brittleness of ionic compounds?

A

Is stress is applied to a crystal of an ionic compounds, then the layers of ionic compounds may slide over each other.
This means ions of the same charge are now side by side and repel each other and therefore the crystals break apart

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19
Q

Why do ionic compounds have high melting points?

A

Ionic solids consist of a giant lattice network of oppositely charged ions. There are many ions in the lattice and the combined electrostatic forces of attraction among all of the ions is large, so a large amount of energy is required to overcome the forces of attraction.

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20
Q

What evidence shows the existence of ions?

A

the ability of an ionic compound to conduct electricity and undergo electrolysis when either molten or in aqueous solution.

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21
Q

Explain the formation of covalent bonds?

A

A covalent bond is formed between two atoms when an atomic orbital containing a single electron from one atom overlaps with an atomic orbital also containing a single electron of another atom.

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22
Q

What is a sigma bond?

A

It is the end-on overlap of two orbitals.

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23
Q

What is a pi bond?

A

A sideways overlap of two p-orbitals leads to the formation of pi bond.

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24
Q

What is a feature of a pi bond?

A

A feature of pi bond is that it cannot form until a sigma bond has been formed. So pi bonds only exist between atoms that are joined by double or triple bonds.

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25
Q

What is dative covalent bonding?

A

Occurs when one atom donates both electrons in a bond.

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26
Q

What is London forces?

A

Weak intermolecular forces arising due to fluctuations of electron density within a nonpolar molecule. These fluctuations may temporarily cause the asymmetric electron distribution: the molecule becomes an instantaneous dipole. This dipole can induce a dipole in another molecule, and so on. The attraction increases with size/shape (points of contact between the molecules) and number of electrons (more fluctuations = more instantaneous/induced dipoles).

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27
Q

What is hydrogen bonding ?

A

A type of intermolecular force between a hydrogen bonded to a more electronegative atom than hydrogen (usually N,O,F) and other atom in a same/different molecule

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28
Q

what is electronegativity?

A

is ability of an atom to attract a bonding pair of electrons.

29
Q

What is polar bond?

A

A polar bond is a covalent bond in which there is a separation of charge between one end and the other.

30
Q

what does the electronegativity of an element depend on?

A
  1. the number of protons in the nucleus
  2. the distance from the nucleus.
  3. the amount of shielding
31
Q

Why does electronegativity increase across a period?

A

Electronegativity increases across a period because the number of chargers on the nucleus increases. That attracts the bonding pair of electrons more strongly.

32
Q

Why does electronegativity fall as you go down the group?

A

As you go down a group, electronegativity decreases because the bonding pair of electrons is increasingly distant from the attraction of the nucleus.

33
Q

What is discrete simple molecule?

A

is an electrically neutral group of two or more atoms held together by chemical bonds.

34
Q

what is covalent bond?

A

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

35
Q

What are london forces?

A

london forces are induced-induced dipole interactions. in any molecule the electrons are moving constantly and randomly. As this happens the electron density can fluctuate and parts of the molecule become more or less negative, i.e. temporary dipoles form.

36
Q

state and explain what factors afffect london forces?

A

number of electrons - more electrons the higher chance temporary dipoles will form. this makes the london forces stronger.
shape of the molecule- long straight alkanes have larger surface area of contact between molecules for london forces to form than compared to spherical shaped branched alkanes and so have stronger london forces.

37
Q

what are permanent dipole-dipole forces?

A

it occurs between polar molcules. stronger than london forces and so the compounds have higher boiling point.

38
Q

What is bond angle around H in hydrogen bond?

A
  1. as there are 2 pairs of electrons around the H atom. these pairs of electrons repel to a position of minimum repulsion.
39
Q

why do alcohols have higher boiling temperatures compared alkanes with similar number of electrons?

A

Alcohols can form hydrogen bond. These require more energy to break theerfore increased boiling temperarures.

40
Q

Explain the density of ice compared to that of water?

why is water more dense?

A

the molecules in ice are arranged in rings of six, held together by hydrogen bonds. The structure creates large areas of open space inside the rings. When ice melts, the ring strcuture is destroyed and the average distance between the molecules decrease,causing an increase in density.

41
Q

Explain the trends in boiling temperatures of the Hydrogen halides?

A

HF has an anamoulsy high boiling point due to hydrogen bonding.
The boiling points then steadily increase from HCl to HI as the number of electrons per molecule increase which in turn results in an increase in London forces.

42
Q

what two conditions must be met in order for a substance to dissolve?

A
  1. the solute partciles must be seperated from each other.
  2. the forces of attraction between the solute and solvent particles must be strong enough to overcome the solvent-solvent forces and the solute-solute forces.
43
Q

explain how ionic solids dissolve 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.
the negative ions are attracted to the hydrogens on the polar water molecule and the positive ions are attracted to the oxygen.
by attracting them they remove them from the lattice.
this process is called HYDRATION.

44
Q

explain why simple alcohols dissolve in water?

A

alcohols contain O-H group and can therefore hydrogen bond with water.
therefore dissolve

45
Q

explain the trend in solubility of alcohols in water?

A

increasing hydrocarbon chain length the london forces predominate between the alcohol molecules.
solubility decreases

46
Q

Explain the insoulibilty of compounds in water?

A

Compounds that cannot form hydrogen bonds with water molecules will be insoluble in water.

47
Q

What is metallic bonding?

A

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

48
Q

What 3 factors affect the strength of metallic bonding?

A
  1. number of protons/strength of nuclear attraction - the more protons the stronger the bond.
  2. number of delocalised electrons per atom - the more electrons the stronger the bond.
  3. size of ion - smaller the ion stronger the bond.
49
Q

Explain the structure of graphite?

A

Each carbon atom is bonded to three others by sigma bonds, forming interlocking hexagonal rings.
planar arragement of carbon atoms in layers.

50
Q

Explain the uses of graphite?

A

Graphite is a fairly good conductor of electricity. The delocalised electrons between layers are free to move under the influence of voltage.
graphite can be used as a solid lubricant since the layers slide over each other easily.

51
Q

Explain the structure and properties of diamond?

A

In diamond each carbon atom forms 4 sigma bonds to 4 other carbon atoms, in a giant three dimensional tetrahederal arrangement. bond angles are 109.5.
diamond is extremely hard due to the very strong C-C bonding throughout the structure. it has very high melting points.

52
Q

explain the strcuture and properties of graphene?

A

one-atom thick layer of graphite.
thinnest material on earth.
200 times stronger than steel.
excellent conductor of both electricity and heat.

53
Q

Uses of graphene?

A

Carbon nanotubes - can be used as a vehicle to deliver drugs to cells.
can be rolled into a ball to produce fullerene molecules.
conductors

54
Q

What are the physical properties in molecular solids?

A

they have low melting and boiling points.
london forces tend to increase with both an increase in the number of electrons per molecule and also with increasing length of molecule.

55
Q

State the main assumptions of the EPR theory?

A

the shape of a molecule/ion is determined by the number of pairs of electron surronding the central atom.
the electron pairs repel one another and arrange themselves in space so that they are as far apart from one another as it possible.
Lone - lone pair repulsion>lone-bond pair>bond-bond

56
Q

what is meant by isoelectronic?

A

contain the same number of electrons.

Same electronic configuration

57
Q

Why is nitrogen much less reactive than either carbon monoxide or hydrogen cyanide?

A

N≡N is stronger than either H-C, C≡N or C≡O

or N≡N is non polar whilst HCN and CO have polar bonds.

58
Q

comment on the difference in boiling temperarures between water and methane?

A

th only force of attraction between methane molcules are london forces, these are very weak and hence methane has low boiling temperature since very little energy is required to overcome these.
Water molcules however have all 3 intermolecular forces. The hydrogen bonds are very strong and therefore it is stronger

59
Q

explain why the ammonia molecule is polar but the boron trochloride molecule is not.

A

in ammonia the three dipoles created by the three polar N-H bonds are acting in the same direction and so they reinforce one another.
In boron trichloride, the three dipoles cancel out since the molecule in symmerical.

60
Q

explain why magnesium is a better conductor of electricity than sodium?

A

The magnesium ion is smaller than the sodium ion and there are two delocalised electrons per ion in magnesium, as opposed to only one per ion in sodium
Hence there are more delocalised electrons per unit volume in magnesium

61
Q

suggest why the electronegativity of fluorine is greater than that of chlorine, despite the fact that the nucleus of a chlorine atom contain more protons?

A

The nucleus of the fluorine atom is closer to the bonding electrons and so the force of attraction for the bonding pair is greater. This factor is more significant than the difference in nuclear charge between fluorine and chlorine.

62
Q

no of bond pairs and lone pairs plus bond angle in trigonal planar?

A
  • 3 bond pairs, 0 lone pairs

- 120

63
Q

no of bond pairs and lone pairs plus bond angle in tetrahedral?

A
  • 4 bond pairs, no lone pairs

- 109.5

64
Q

no of bond pairs and lone pairs plus bond angle in trigonal bipyramidal?

A

-5 bond pairs , 0 lone pairs

90 and 120

65
Q

no of bond pairs and lone pairs plus bond angle in V shaped?

A
  • 2 bond pairs, 2 lone pairs

- 104.5

66
Q

no of bond pairs and lone pairs plus bond angle in trigonal pyramidal?

A
  • 3 bond pairs, 1 lone pair

- 107

67
Q

no of bond pairs and lone pairs plus bond angle in octahederal?

A
  • 6 bond pairs,0 lone pairs

- 90 and 180

68
Q

no of bond pairs and lone pairs plus bond angle in linear?

A
  • -2 bond pairs

- 180

the underlying truth of chemistry explored by me (11 decks)

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