CH2: Structure Flashcards
What is electronegativity?
Power of an atom to attract the pair of electrons in a covalent bond towards itself
How does nuclear charge effect electronegativity?
- attraction between protons in nucleus and electrons on outer shell
- this attraction increases when no. Protons in nucleus increases
- increased NUCLEAR CHARGE , INCREASED ELECTRONEGATIVITY
How does the atomic radius effect electronegativity?
Electrons closer to nucleus - more strongly attracted to positive nucleus
an increased atomic radius results in a decreased electronegativity
How does shielding effect electronegativity?
- 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
Trend of electronegativity down the group?
DECREASE in electronegativity
1) nuclear charge increases
2) HOWEVER , the elements have more SHELLS as you go down the group —> INCREASE SHIELDING/RADIUS
Trend in electronegativity across a period ?
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
When is the covalent bond NON POLAR?
When tow atoms have same electronegativity
How does the difference in electronegativities determine the bond formed?
- 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
Which part of the atom will be positive and which art will be negative if the atom is POLAR?
Less electronegative atom - delta positive
More electronegative atom- delta negative
How to determine the polarity of a molecule?
- the polarity of each bond
- how the bonds are arranged in the molecule (if symmetrical - not polar)
What is metallic bonding?
electrostatic forces of attraction between + metal ions and delocalised electrons
Properties of GRAPHITE ?
CONDUCT ELECTRICITY : 3 covalent bonds per atom, so 4th electron per atom is delocalised
Cannot conduct electricity from one layer to the next as energy gap between layers is too large for easy electron transfer
SOFT AND SLIPPERY : weak intermolecular forces between layers - can slide
Structure of DIAMOND?
4 Covalent bonds per atom —> TETRAHEDRON
- NO INTERMOLECULAR FROCES
Structure of GRAPHENE?
Single layer of graphite - one atom thick
3 covelnt bonds per atom - 4th electron delocalised - conduct electricity
High tensile strength due to strong covalent bonds
Solubility of all structures?
GIANT IONIC : soluble
GIANT METALLIC : insoluble
SIMPLE COVALENT : insoluble unless polar
GIANT COVALENT : insoluble
What are van der waal forces?
- 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
What are permanent dipole forces?
- between polar molecules (have both. Negative and positive charged end)
- oppositely charged ends attract
What makes London forces stronger?
- more electrons so stronger London forces
Trend in BP in group 6 hydrides ?
-BP increase from H2S to H2Te due to increased no. Electrons = more London forces
Water doesn’t follow this pattern - BP larger than hydrides due to Hydrogen Bonds in water but not in other hydrides
BP trend in group 7 hydrides?
HF molecules form hydrogen bonds with each other - highest BP
From HCl to HI, BP increases due to more electrons = more London forces
When do hydirgen bonds occur?
When hydrogen is covakently bonded to FLUORINE, NITROGEN AND OXYGEN
- form bonds with lone pairs of electrons on fluorine, oxygen and nitrogen
Why is ice less dense than water?
In ice , water molecules are arranged in a lattice/hexagon
- arranged so more space between molecules - hydrogen bonds longer than covalent bonds- less dense
When melted , some hydrogen bonds are broken/lattice breaks down - molecules get closer/more tightly packed together
When will a substance dissolve?
Strength of new bonds formed is same as/greater than strength of bonds broken
2 types of solvent?
Polar solvents - e.g water - solute and solvent can bond by H bonding/dipole-dipole interactions
(Polar substances dissolve in polar solvents)
Non polar solvents - e.g hexane
- hexane molecules bond together by London forces
(Non polar substances dissolve in non polar solvents)
How do ionic substances dissolve in polar solvents ?
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
How do alcohols dissolve in polar solvents ?
Bc polar O-H bond attracted to O-H bonds in water - form hydrogen bonds as there are lone pairs on oxygen
Carbon chain isn’t attracted to water so more C atoms there are , the less soluble alcohol will be
Why don’t all molecules with polar bonds dissolve in water?
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
What substances tend to dissolve best in non polar solvents?
Non polar substance - form similar bonds (London forces)
Size of O-H-O angle in water ?
180 degrees
