Topic 2: Bonding Flashcards
Define Ionic Bonding
the strong electrostatic attraction between the oppositely charged ions
what does charge density depend on?
ionic charge and ionic radius
what is proof for the existence of ions
migration of ions and physical properties
migration of ions example
Copper (II) Chromate
copper 2+ ions migrate to negative electrodes (blue)
chromate ions migrate to the positive electrode (yellow)
define covalent bonding
the strong electrostatic attraction between the shared pair of electrons and the nuclei of the atoms bonded
what is dative covalent bonding
instead of one electron coming from each atom, the electrons come from a single atom
below which period in the periodic table can atoms expand their octet
below period 2
define bond energy
measure of how strong a bond is
what does a double bond consist of?
a sigma bond and pi bond between 2 atoms
angle in a tetrahedral molecule
109.5
angle in a pyramidal molecule
107
angle in a bent molecule
104.5
define electronegativity
the ability of an atom to attract the bonding electrons in a covalent bond
what are London forces
when a temporary dipole is attracted to an induced dipole
define metallic bonding
strong electrostatic attraction between metal ions and the delocalised electrons
tetrachloromethane, CCl₄, contains polar bonds but is not a polar molecule, why?
the centres of charges are in the same place therefor it is non polar
explain why the melting temp of magnesium is much higher than that of sodium
magnesium ion has a larger charge density
magnesium ions are smaller than sodium ions
magnesium ions have more delocalised electrons
magnesium ions have a greater attraction for the delocalised sea of electrons
the electrical conductivity of pure silicon is very low. explain why this is in terms of the bonding.
silicon’s outer electrons are fixed in covalent bonds
therefore silicons electrons are not free to move
explain why sulphur forms an expanded octet but oxygen does not
to expand the octet an electron as to be promoted to a higher energy level
for sulfur the next energy level is fairly close, for oxygen the energy jump is too large
describe the shape of TeF₆
octahedral. 90 degree angle
describe the shape of AsF₅
trigonal bipyramidal. 120 and 90 degrees
describe the shape of PH₃
trigonal pyramidal. 107.5 degrees
state and explain whether you would expect magnesium fluoride to have a higher or lower melting point than magnesium oxide
magnesium fluoride has lower
fluoride has a lower charge than oxide
less energy needed to overcome/ break bonds
suggest why copper (II) chromate ions dissolve in water
ions are hydrated
state and explain the trend in ionic radius from K+ to Ti4+
ionic radius decreases
ions are isoelectronic
nuclear charge increases
attraction between electrons and nucleus increases
aluminium fluoride (AlF₃) melts at a much higher temp than either aluminium bromide or aluminium chloride, but none of the compounds conduct electricity when molten. compare the structure and bonding present in the three compounds. (4)
AlF₃ bonding is covalent
structure is giant
other bonding is covalent
structure is simple molecular
explain why hydrogen halides are gases at room temp
bp lower than room temp
weak intermolecular forces
little energy needed to overcome the forces between the molecules
explain why less energy is needed to break the covalent bond in hydrogen iodide than in hydrogen bromide (3)
iodine atoms are larger than bromine atoms
bond is longer in HI than HBr
attraction between nuclei and shared pair of electrons is weaker
state and explain the trend in bond polarity in the hydrogen halides as atomic number of the halogen increases
bond polarity decreases with atomic number
halogen atom becomes less electronegative
smaller difference in electronegativity between H and halogen
explain why the F-P-F bond angle in PF₃ (107) is different to the F-Si-F bond angle in SiF₄ (109.5)
electron pair repulsion pushing bonds closer together
lone pairs repelling more than bonding pairs (PF₃)
Explain why CH₄ molecules are tetrahedral and state the H-C-H bond angle
electron pairs repel
shape provides minimum repulsion
angle is 109.5
Explain why the H-O-H bond angle in water is smaller than the H-N-H bond angle in NH₃
water has one more lone pair of electrons than NH₃
lone pairs repel more than bonding pairs
bonding pairs are pushed closer
explain why HF has a higher melting point than HCl
HF forms hydrogen bonds
HCl only forms London forces
hydrogen bonds are stronger
explain why HF dissolves better in water than hexane
both involve breaking hydrogen bonds (between HF molecules)
hydrogen bonds are reformed when HF dissolves in water
only London forces formed when HF dissolves in hexane
describe how HF molecules form London forces
random movement of electrons temporary dipole forms induced dipole (in adjacent molecules)
explain why CH₃CH₂CH₂CH₃ has a higher boiling point than (CH₃)₃CH
branched alkane/ (CH₃)₃CH has weaker London forces
molecules have lower SA/ cant be as close together
why are alkanes insoluble in water
they cant form hydrogen bonds
Explain why methanethiol (CH₃SH) cant form hydrogen bonds
sulphur is less electronegative than oxygen
S-H bond is not polar enough (to attract S: from another molecule)
state and explain which of methanol and methane is more volatile
methane
cant form hydrogen bonds
hydrogen bonds are stronger (than London forces)
Explain why methanol molecules are polar and describe a lab test to prove it
(C-O and) O-H bonds polar
not symmetrical
test: charged object near jet of methanol
result: jet deflected
describe the hydration of the chloride ions formed when HCl dissolves in water
partial positive H atom of water (1) is attracted to (full) negative charge of chlorine ions
suggest why HCl dissolves well in methylbenzene (C₇H₈)
similar intermolecular forces
London forces
why is ice less dense than water?
more open/more space between molecules (making it less dense) due to (3 Dimensional) lattice/ ring structure in ice
hydrogen bonds longer than the covalent bonds
Aluminium fluoride and aluminium chloride are both crystalline solids at room temperature. Aluminium fluoride sublimes to form a gas at 1291qC (1564 K), whilst aluminium chloride sublimes at 178qC (451 K).
Use the Pauling electronegativity values in the Data Booklet to explain these differences in sublimation temperature.
aluminium and chlorine electronegativity difference 1.5 AND aluminium and fluorine electronegativity difference 2.5 aluminium chloride (mostly) covalent / (small) molecule aluminium fluoride (bonds) more polar aluminium chloride molecular so weak(er) intermolecular forces / London forces aluminium fluoride is a giant structure/ strong electrostatic forces of attraction between the ions more energy needed to break the stronger bonds to cause sublimation in aluminium fluoride
