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IB Chemistry HL > UNIT 4 - Chemical Bonding And Structure > Flashcards

UNIT 4 - Chemical Bonding And Structure Flashcards

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1
Q

State the angle with which molecules with three electron domains will position and name it and list an exception to it.

A
  • 120 degrees
  • triangular planar

exception: if one or more electron domain is a lone pair

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2
Q

State the types of orbital overlapping in sigma bond and pi bond.

A

sigma - axial overlapping
pi - sideways overlapping

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3
Q

State the angle with which molecules with three electron domains (2 bonding and one lone) will position and name it.

A

bent, v-shaped
117 degrees

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4
Q

State how does presence of polarity in a bond affect its character.

A

introduced ionic characteristics into it so the more polar the bond the more ionic-like the compound behaves like

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5
Q

State the angle with which molecules with four electron domains will position and name it.

A

tetrahedral shape
109.5 degrees

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6
Q

State and explain the physical properties of ionic compounds.

A
  • melting and boiling point - high => strong electrostatic attraction between the lattice ions and thus a lot of energy to break them
  • solids at room temp
  • volatility - low => they are crystalline solids
  • solubility - very well in polar solvents like water => partial charges in water molecules are attracted to ions of opposite charge in the lattice which causes ions to dislodge and become surrounded by water molecules (hydrated)
  • conductivity - no in the solid state, yes in aqueous solutions and melted => must have ions that move like jagger and carry charge
  • brittleness - high
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7
Q

Define a net dipole moment.

A

Either the molecules contain bonds of different polarity, or its bond is not symmetrically arranged the dipoles will not cancel out and thus the molecule will be polar

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8
Q

State the angle with which molecules with two electron domains will position and name it.

A

180 degrees
linear shape

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9
Q

Define delocalized electrons, their function and state where they are present.

A

Electrons with a tendency to be shared between more than one bonding position, are free from the constraints of single bonding position.

give greater stability to the molecule or ion

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10
Q

Define a polar bond.

A

A bond between two atoms of different electronegativity values in which one atom exerts a stronger pulling power than the other one and thus the bond is unsymmetrical in terms of electron distribution.

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11
Q

Define when resonance occurs in molecular structure.

A

When more than one valid Lewis structure can be drawn for a particular molecule.

The true structure is an average of resonances known as resonance hybrid.

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12
Q

Distinguish between cation and anion.

A

atom loses electrons and forms a positive ion = cation

atom gains electron and forms a negative ion = anion

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13
Q

What does a circle inside a hexagon represent for organic molecules.

A

Delocalized pi electrons spread equally through the ring, rather than being confined double bonds.

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14
Q

Define a coordinate bond and state its alternative name.

A

dative bond

A covalent bond in which both the shared electrons are provided by one atom.

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15
Q

Define a giant molecular structure and list its two other possible names.

A

Network covalent, macromolecular structure

Crystalline lattice in which the atoms are linked together by covalent bonds, so no finite size

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16
Q

State the order in which the electron domain repulsion decreases.

A

lone pair-lone pair > lone pair-bonding pair > bonding pair-bonding pair

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17
Q

Define allotropes and list an example.

A

Different structural modifications of an element in the same physical state.

Usually have different properties

e.g. oxygen (O2) and ozone (O3)

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18
Q

State and define the structure of ionic compounds and what does coordination number mean in them.

A

lattice structure

three-dimensional, crystalline structure where ions surround themselves with ions of opposite structure, can grow inifinitely as it is just an expresssion of repeated “unit” - formula unit

coordination number describes the number of ions that surround a given ion in the lattice (6 in NaCl)

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19
Q

List the carbon allotropes.

A
  • graphite
  • diamond
  • fullerene
  • graphene
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20
Q

Define a covalent bond.

A

An electrostatic attraction between a pair of electrons and positively charged nuclei.

Involves sharing an electron pair between two atoms aiming to gain electrons.

Can be single (sigma), double (sigma and pi) or triple (sigma and 2 pi).

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21
Q

Describe graphite taking into consideration its structure, electrical conductivity, thermal conductivity, appearance, special properties and use.

A

structure:
each carbon is sp2 (covalently bonded to 3 others), forms hexagons in parallel layers with bond angles 120 degrees.
Layers held by dispersion forces so they can slide over each other.

conductivity:
great because of the delocalized p electrons between layers

thermal:
not really unless in direct parallelism to the crystal layers

app:
non-lustrous, grey, crystalline solid

sp:
slipping layers over each other, soft, brittle, very high melting point, most stable

use:
pencils electrode rods in electrolysis

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22
Q

List the two characteristics of the covalent bond, define and state how they relate to the strength of the bond.

A
  • bond length - measure of distance between two bonded nuclei.
  • bond strength (bond enthalpy) - measure of the energy required to break the bond.

the shorter the stronger the bond

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23
Q

Describe diamond taking into consideration its structure, electrical conductivity, thermal conductivity, appearance, special properties and use.

A

structure:
each carbon sp3 , so to 4 others, angles 109.5 degrees

e conductivity:
non-conductor, all electrons bonded

thermal:
very efficient, better than metals

app:
highly transparent, lustrous, crystal

sp:
hardest known natural substance, brittle, ultra high melting point

use:
jewelry, grinding and cutting glass

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24
Q

Define the term electron domain and what does it determine.

A

All electron locations on the valence shell occupied by lone pairs, single, double, triple bonded pairs whatever.

Determines the geometrical arrangement and thus the shape of a molecule.

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25
Q

Define solubility in terms of physical properties.

A

Ease with which a solid (the solute) becomes dispersed through a liquid (the solvent) to form a solution.

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26
Q

Describe fullerene taking into consideration its structure, electrical conductivity, thermal conductivity, appearance, special properties and use.

A

structure:
sp2 hybridization, bonded in a sphere of 60 carbon atoms, a cage, fixed formula.

electrical conductivity:
semiconductor in normal temp

their conductivity:
very low conductivity

app:
yellow crystalline solid, soluble in benzene

sp:
very light and strong, low melting point

use:
medical industrial devices for binding specific target molecules, lubricants, nanotubes

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27
Q

Define the ionic bond.

A

An electrostatic attraction between oppositely charged ions.

28
Q

Describe graphene taking into consideration its structure, electrical conductivity, thermal conductivity, appearance, special properties and use.

A

structure:
covalently bonded to 3 other carbons as in graphite but just single layer, described as honeycomb

electrical conductivity:
very good, one delocalized electron per atom

thermo conductivity:
best known, better than diamond even
app: almost completely transparent

sp:
ultra strong ultra thin, very flexible, very high melting point

use:
TEM, tech in general, electronic devices

29
Q

Expand the abbrev VSEPR and state the simple notion on which it is based.

A

Valence Shell Electron Pair Repulsion theory

Electron pairs in the same valence shell carry the same charge they repel each other and so spread themselves as far apart as possible.

30
Q

State the structure formed by silicon dioxide and its properties.

A

Giant covalent structure based on tetrahedral arrangement.

  • strong
  • insoluble in water
  • high melting point
  • non-conductor of electricity
31
Q

Define volatility in terms of physical properties.

A

Tendency of substance to vaporize.

32
Q

List one exception when even with electronegativity difference a molecule is non-polar.

A

PH3

33
Q

State compounds/atoms that are prone to extended or incomplete octet formation.

A
  • B(BF3)
  • SF6
  • AlCl3
  • PCl5
34
Q

Define intermolecular forces.

A

Forces existing between molecules
which determine the physical properties of a substance.

35
Q

State the electronegativity difference that in most cases allows us to assume a compound is ionic.

A

> 1.8

36
Q

List the forces counting as van der Waals forces.

A
  • London dispersion forces
  • dipole-dipole forces
  • dipole - instantaneous dipole
37
Q

Define a dipole.

A

Indicator of the fact that a bond has two separated opposite electric charges.

38
Q

State what happens in dipole - instantaneous dipole forces.

A

Dipole of uneven charges (induced dipole) is created very quickly and randomly in a molecule because of an influence of another molecule.

39
Q

Define lattice energy.

A

Measure of the strength of the attraction between ions within the lattice (greater for small, high charge).

40
Q

State how London dispersion forces occur.

A

Weak forces of attraction will occur between opposite ends of two temporary dipoles in a molecule.

41
Q

State when and why does strength of London dispersion forces increase.

A

With increasing molecular size,
because the greater no of electrons within a molecule the greater probability of temporary dipoles developing.

42
Q

State the properties of molecules with London dispersion forces.

A

Exist also in non-polars remember.

Low melting and boiling points because they are weak.

Many are gases at room temp because of that
are responsible for the fact that non-polar molecules can be condensed to form liquids and sometimes solids.

43
Q

State where and how dipole-dipole attractions occur.

A

In polar molecules because they have permanent separation of charge within their bonds (the difference in electronegativity) they do possess a permanent dipole and thus when neighboring molecules have opposite charges they attract.

44
Q

Define a permanent dipole.

A

One end of molecule is electron deficient with partial positive charge while the other is electron rich with partial negative charge.

45
Q

State on what does the strength of dipole-dipole force depend on.

A

Distance and relative orientation of the dipoles.

46
Q

Define a hydrogen bond.

A

An intermolecular attraction between molecules containing hydrogen covalently bonded to a very electronegative atom (FON).

Because of small size of hydrogen and no other shielding it exerts a strong attractive force on lone pair in the electronegative atom of a neighboring molecule.

47
Q

Which intermolecular force is the strongest and which is the weakest.

A

strongest - hydrogen
weakest - London

48
Q

State the forces responsible for the ice properties.

A

hydrogen

49
Q

Explain why most covalent structures are liquid or gases at room temperature.

A

They have also lower boiling and melting points than ionic compounds because the forces to overcome to separate the molecules are relatively weak intermolecular forces so way easier to break than the electrostatic attractions in the ionic lattice.

50
Q

Explain the macromolecular or giant covalent structures melting and boiling point and list an example of such.

A

very high, often as high as those of ionic compounds

covalent bonds must be broken in these compounds for the change of state to occur

e.g. diamond

51
Q

State why non-polar substances are generally able to dissolve in non-polar solvents.

A

Because of the formation of London dispersion forces between solute and the solvent.

52
Q

State why polar substances are generally able to dissolve in polar solvents.

A

Because of dipole interactions and hydrogen bonding between solute and solvent.

53
Q
A
54
Q

State where and why the solubility of polar compounds is reduced.

A

Is reduced in plarger molecules where the polar bond is only a small part of the total structure, non-polar parts unable to associate with water reduce its solubility.

54
Q

Define a metallic bond.

A

Electrostatic attraction between the lattice of cations and the delocalized electrons.

54
Q

List the factors that determine the strength of the metallic bond.

A
  • no of delocalized electrons
  • charge of the cation
  • radius of the cation

e.g. the greater the no of delocalized e and the smaller the cation the greater the binding force between them.

54
Q

Describe the physical properties of metals taking into consideration its bonding type, electrical conductivity, thermal conductivity, appearance, malleability, ductility, melting points and use.

A

bonding: metallic

  • good electrical conductivity:
    delocalized electrons that are very mobile and thus can move through the metal structure in response to applied voltage; electrical circuits with the use of copper
  • good thermal conductivity:
    delocalized electrons and closed packed ions enable efficient transfer.
  • malleable and ductile:
    movement of delocalized electrons is non-directional and random so the metallic bond remains intact while the conformation changes under applied pressure
  • high melting points:
    strong metallic bonds
  • shiny, lustrous:
    delocalized electrons in crystal structure reflect light.
54
Q

Define alloys.

A

Solid solutions usually containing more than one metal and held together by metallic bonding, usually with different properties, chemically more stable and more resistant to corrosion than their component elements.

Possible because of non-directional nature of delocalized electrons and the fact that lattice can accommodate ions of different structure.

54
Q

State exceptions to the octet rule.

A
  • Small atoms (be, B) form stable molecules with fewer than an octet of electrons.
  • Atoms of elements in Period 3 and below may expand their octet by using d orbitals in their valence shell.
54
Q

State when, according to its molecular geometry, a molecule is non-polar.

A

If there is no net dipole: no lone pairs and all atoms attached to the central atom are the same.

54
Q

State the geometry shapes and angle of molecules with six electron domains.

A

90 degrees
octahedral shape - 0 lone pairs
square pyramidal - 1 lp
square planar - 2 lp

54
Q

Define a sigma bond, types of orbitals it is formed from and how.

A

A covalent bond formed by axial overlapping of sp and/or hybrid orbitals.

54
Q

State the factors ensuring a most stable Lewis structure.

A
  • the lowest formal charge
  • negative values for formal charge on the more electronegative atoms.
54
Q

Define a free radical and list an example.

A

reactive species that contains an unpaired electron

e.g. NO

54
Q

State the formula for formal charge.

A

FC= V (valence e) - (1/2B (no of bonding electrons) +L (no of lone e pairs))

54
Q

Define hybridization.

A

Process where atomic orbitals within an atom mix to produce hybrid orbitals of intermediate energy, the atom is able to form stronger covalent bonds using these orbitals.

54
Q

Associate the hybridizations with the types of molecular geometry.

A

sp - linear
sp2 - triangular planar
sp3 - tetrahedral

IB Chemistry HL (13 decks)

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