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Chemistry A- Level > CH2: Structure > Flashcards

CH2: Structure Flashcards

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1
Q

What is electronegativity?

A

Power of an atom to attract the pair of electrons in a covalent bond towards itself

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2
Q

How does nuclear charge effect electronegativity?

A
  • attraction between protons in nucleus and electrons on outer shell
  • this attraction increases when no. Protons in nucleus increases
  • increased NUCLEAR CHARGE , INCREASED ELECTRONEGATIVITY
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3
Q

How does the atomic radius effect electronegativity?

A

Electrons closer to nucleus - more strongly attracted to positive nucleus

an increased atomic radius results in a decreased electronegativity

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4
Q

How does shielding effect electronegativity?

A
  • filled shells can Shield the effect of nuclear charge —> outer electrons are less attracted to nucleus
    an increased number of inner shells and subshells will result in a decreased electronegativity
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5
Q

Trend of electronegativity down the group?

A

DECREASE in electronegativity
1) nuclear charge increases
2) HOWEVER , the elements have more SHELLS as you go down the group —> INCREASE SHIELDING/RADIUS

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6
Q

Trend in electronegativity across a period ?

A

INCREASES
- nuclear charge increases
- shielding is constant (all have same no. Shells/sub shells))
- so as the attraction between nucleus and outer electrons is strong , ATOMIC RADII DECREASES

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7
Q

When is the covalent bond NON POLAR?

A

When tow atoms have same electronegativity

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8
Q

How does the difference in electronegativities determine the bond formed?

A
  • when difference between electronegativities is more than 1.7 - IONIC BOND and ions formed
  • when difference is 0.3 to 1.7 - COVALENT BOND and is POLAR
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9
Q

Which part of the atom will be positive and which art will be negative if the atom is POLAR?

A

Less electronegative atom - delta positive
More electronegative atom- delta negative

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10
Q

How to determine the polarity of a molecule?

A
  • the polarity of each bond
  • how the bonds are arranged in the molecule (if symmetrical - not polar)
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11
Q

What is metallic bonding?

A
  • lattice structure of positive charge in a ‘sea’ of delocalised electrons held by electrostatic forces between opposite charges
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12
Q

Properties of GRAPHITE ?

A

CONDUCT ELECTRICITY : each atom bonded to 3 others, so 1 DELOCALISED

SOFT AND SLIPPERY : weak intermolecular forces between layers - can slide

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13
Q

Structure of DIAMOND?

A

each carbon bonds with 4 others —> TETRAHEDRON

  • NO INTERMOLECULAR FROCES
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14
Q

Structure of GRAPHENE?

A

Single layer of graphite - one atom thick

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15
Q

Solubility of all structures?

A

GIANT IONIC : soluble
GIANT METALLIC : insoluble
SIMPLE COVALENT : insoluble unless polar
GIANT COVALENT : insoluble

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16
Q

What are van der waal forces?

A
  • exist between all atoms/molecules
    Weak intermolecular forces arising due to fluctuations in electron density in a non polar molecule
  • electron charge cloud in non polar molecules move constantly —> causes temporary dipole
  • temporary dipole can induce a dipole on neighbouring molecules
  • δ+ end of the dipole in one molecule and the δ- end of the dipole in a neighbouring molecule are attracted towards each other
17
Q

What are permanent dipole forces?

A
  • between polar molecules (have both. Negative and positive charged end)
  • oppositely charged ends attract
18
Q

What makes London forces stronger?

A
  • more electrons so stronger London forces
19
Q

Trend in BP in group 6 hydrides ?

A
  • enthalpy of vaporisation increase from H2S to H2Te due to increased no. Electrons = larger London forces
    Water doesn’t follow this pattern - enthalpy change of vaporisation larger than hydrides due to Hydrogen Bonds in water but not in other hydrides
20
Q

BP trend in group 7 hydrides?

A

HF molecules form hydrogen bonds with each other - highest BP
From HCl to HI, BP increases due to more electrons = strength of London forces increases

21
Q

When do hydirgen bonds occur?

A

When hydrogen is covakently bonded to FLUORINE, NITROGEN AND OXYGEN
- form bonds with lone pairs of electrons on fluorine, oxygen and nitrogen

22
Q

Why is ice less dense than water?

A

In ice , water molecules are arranged so there is max number of H bonds - lattice ‘wastes’ space
- arranged so more space between molecules - longer bond lengths - less dense

When melted , some hydrogen bonds are broken/lattice breaks down - molecules ‘fill’ spaces

23
Q

When will a substance dissolve?

A

Strength of new bonds formed is same as/greater than strength of bonds broken

24
Q

2 types of solvent?

A

Polar solvents - e.g water - molecules bond by H bonds / can only form London forces/permanent dipole forces

Non polar solvents - e.g hexane
- hexane molecules bond together by London forces

25
Q

How do ionic substances dissolve in polar solvents ?

A

Ions attracted to oppositely charged ends of water molecules
- ions pulled away from ionic lattice by water molecules , which surrounds ions - HYDRATION

They will dissolve if enthalpy change of hydration is greater than/closer to energy needed to break up lattice

Some ionic substances don’t dissolve bc bonding of ions too strong

26
Q

How do alcohols dissolve in polar solvents ?

A

Bc polar O-H bond attracted to O-H bonds in water - form hydrogen bonds to lone pairs on oxygen

Carbon chain isn’t attracted to water so more C atoms there are , the less soluble alcohol will be

27
Q

Why don’t all molecules with polar bonds dissolve in water?

A

Halogenoalkanes have polar bonds but DIPOLES aren’t strong enough to form H BONDS with water
- hydrogen bonding in water is stronger than bonds that would be formed with halogenoalkanes

However form permanent dipole forces so dissolve in polar solvents that form permanent dipole forces

28
Q

What substances tend to dissolve best in non polar solvents?

A

Non polar substance - form similar bonds (London forces)

29
Q

Size of O-H-O angle in water ?

A

180 degrees

Chemistry A- Level (23 decks)

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