Module 2 - [ch5,6] (atoms) Module 3 - [ch7] (periodicity) Flashcards
electrons and bonding shapes of molecules and intermolecular forces periodicity
Why does first ionisation energy decrease going down a group [3]
atomic radius increases,
increased distance between outermost electron and nucleus
more electron shells
more electron shielding
less nuclear attraction
How are shells filled with electrons?
orbitals are filled in order of energy (lowest energy shells are filled first)
electrons are assigned individually before being paired
what is an atomic orbital
a region around the nucleus holding max two electrons with opposite spin
(cloud of negative charge)
it’s impossible to know where exactly the electron is
ammonium ion (charge)
NH4+
what is the shape and bond angle of an NH4+ ion
tetrahedral , 109.5
Discuss the conductivity of ionic compounds [2]
when solid ions are in fixed positions and are not free to carry charge
when molten or in aqueous solution ions are mobile and can carry charge
define first ionisation energy [3]
energy required to remove one electron [1] from each atom in 1 mole [1] of gaseous atoms [1] to form 1 mole of gaseous 1+ ions
why is second ionisation energy generally greater than the first
since an electron has been removed the remaining electrons experience less repulsion and are pulled closer to the nucleus
experience more nuclear attraction
harder to remove
proton : electron ratio is greater
define isotopes [1]
atoms of an element with the same number of protons but different number of neutrons [1]
what would be evidence for shells in an atom?
jump in first ionisation energy
suggests the electron is much harder to remove since is much closer to the nucleus
Explain why the first ionisation energies show a general increase across
Period 2 [2]
increasing atomic number ( number of protons)
more nuclear attraction
similar shielding across period since same shell
Explain why the first ionisation energy of B is less than that of Be. [2]
In B, an electron is being removed from a higher energy level
An s electron lost in Be and a p electron lost in B
What determines which block an element is in?
the highest energy level shell being filled
what is meant by ionic bonding [1]
electrostatic force of attraction between oppositely charged ions [1]
predict the bond angle of an F2O molecule, explain your choice [3]
2bp , 2lp
bond angle 104.5
bent (non-linear)
lone pairs repel more than bonded pairs
Describe and explain two anomalous properies of water which results from
hydrogen bonding [4]
when solid, water is less dense than when liquid
ice has an open lattice
relatively high mp and bp
hydrogen bonds are strong, require lots of energy to overcome
what is metallic bonding
giant metallic lattice
electrostatic attraction between fixed positive ions and delocalised electrons
describe the structure and bonding shown by cl, how does this explain it’s difference in melting point compared to Mg ?
simple molecular lattice
Cl has london forces between molecules
london forces weaker than metallic bonds hence a lower melting point than Mg
Explain, in terms of bonding and structure, the properties of graphite [6]
graphite is maleable [1] , has a very high melting point [1] and a good conductor [1]
maleable - layers can slide over one another, weak london forces between layers [1]
high mp - strong covalent bonds require energy to overcome (giant covalent lattice) [1]
good conductor - delocalised electrons free to carry charge [1]
shape and bond angle of an ammonia molecule
NH3
3bp 1lp
107
pyramidal
lone pairs repel more than bonded pairs
diagram of hydrogen bonding?
H+ - O:~ —— H+ - O: ~
Suggest why H2S has a much lower boiling point than H2O
no hydrogen bonding,
weaker IM forces, less energy required to overcome forces
why is chlorine is a gas at room temperature but carbon does not boil until well over 4500 °C.
Explain this difference, in terms of bonding and structure.
to boil Cl2 weak london forces have to be broken
to boil carbon covalent bonds need to be broken
covalent bonds are much stronger than london forces, more energy required
what is meant by hydrogen bonding
electrostatic attraction between a hydrogen atom in one polar molecule (e.g. water) and a small electronegative atom (O,N,F) in another molecule
define electronegativity
ability of an atom to attract the pair of electrons in a covalent bond
explain a polar bond in terms of electronegativity
one element is more electronegative than the other
attract the pair of electrons better,
imbalance of charge
?????
why does boiling point increase from Na to Al
From Na → Al, no of delocalised electrons increases
charge on positive ion increases/
ionic size decreases/
charge density increases
attraction between + ions and electrons increases/
metallic bonding gets stronger
Name the shape and bond angle of an NCl3 molecule
why does it have this shape and bond angle
pyramidal
107
electron pairs repel to get as far apart as possible
lone pairs repel more than bonded pairs
3bp, 1lp surround central N atom
base + acid –> salt + water
why does electrical conductivity of solution decrease as acid is added ?
as acid is added, reactions take place using up ions which were carrying charge
The O–H bonds in water and the N–H bonds in ammonia have dipoles
why do they have dipoles? [1]
one of the elements is more electronegative than the other [1]
more electronegative one attracts the pair of electrons closer
Describe and explain the density of ice compared with water [2]
ice is less dense than water
hydrogen bonds cause an open lattice structure in ice [2]
What is the difference between a covalent bond and a dative covalent bond? [1]
dative covalent, bonded pair comes from same atom [1]
why is chlorine is a
stronger oxidising agent than iodine.
Cl has fewer shells
electron will be more strongly attracted
Explain why a water molecule has a different shape from a carbon dioxide
molecule. [2]
electron pairs repel to get as far apart as possible
water - non-linear 104.5
co2 - 180 linear
Oxygen in water surrounded by 4 areas of electron density/2 bonds and 2 lone pairs
Carbon in CO2 surrounded by 2 regions of electron density/2 double bonds
Explain why water has polar molecules but carbon dioxide has non-polar molecules [2]
co2 is symmetrical whereas h20 isn’t
in co2 dipoles cancel out
silver
Ag+
zinc
Zn2+
aluminium
Al3+
ammonium
NH4+
carbonate
C03 2-
sulfate
SO4 2-
hydroxide
OH-
nitrate
NO3-
how does solubility work
ionic lattice must be broken down
water molecules must surround ions and attract them
solubility decreases as ionic charge increases since attraction may be too strong to overcome
properties of ionic compounds
high mp and bp
soluble in polar solvents
conductors in molten or aq state
what is a covalent bond really
orbital overlap
what does it mean to say a covalent bond is localised and an ionic bond is not
in an ionic bond the ion attracts in all directions
in a covalent bond the attraction acts solely between the shared pair of electrons and the nuclei
define average bond enthalpy
measurement of covalent bond strength
measure of average energy needed to break the bond
what is electron-pair repulsion theory
the electron pairs surrounding a central atom determine the shape of the molecule
electron pairs will repel each other to get as far apart as possible to minimise repulsion
lone pairs repel each other more strongly than bonded pairs
4bp
tetrahedral 109.5
3bp
trigonal planar 120
2bp
linear 180
1 more lone pair does what to the bond angle
decreases it by 2.5
3bp 1lp
trigonal pyramidal 107
2bp 2lp
non linear (bent) 104.5
6bp
octahedral 90 SF6
how is electronegativity measured
pauling electronegativity scale
increasing from bottom left to top right
rough values of pauling electronegativity
0 - covalent
<= 1.8 - polar covalent
> 1.8 - ionic
how is the electron pair shared in a non-polar bond
equally between the bonded atoms
both have the same / similar electronegativity
what does the strength of london forces increase with
number of electrons
which of the IM forces are the strongest and weakest
hydrogen bonding
london forces
what is soluble in what?
polar subtances in polar solvents
non-polar substances in non-polar solvents
3 factors affecting electronegativity
atomic radius
nuclear charge (number of protons)
shielding
why is the 4s shell filled first?
what are the exceptions to this and why?
lower energy than 3d
chromium and copper
half/full 3d subshell more stable
properties of simple molecular substances
low bp
weak intermolecular forces
strong covalent bonds (intramolecular)
not affected when boiled
relatively small molecules
doesn’t conduct
define covalent bonding
electrostatic attraction between nuclei and a shared pair of electrons
why are ionic compounds soluble in polar solvents
polar solvents such a water have a polar bond
the delta positive and delta negative charges are able to attract the charged ions
what is a lone pair
a pair of electrons on the outer shell not involved in bonding
what does expansion of the octet mean
when a bonded atom has more than 8 atoms in its outer shell
what does expansion of the octet mean
when a bonded atom has more than 8 atoms in its outer shell
describe the bonding in simple molecular substances
atoms within the same molecule are bonded together by strong covalent bonds
different molecules are held together by weak intermolecular forces
what are three properties of giant covalent structures
high melting and boiling points
non-conductors except graphite
insoluble in both polar and non-polar solvents
what does it mean to say a bond is non-polar
the electrons in the bond are evenly distributed
what is meant by intermolecular force
attractive force between neighbouring molecules
how are the elements arranged in the periodic table
increasing atomic number
what is meant by periodicity
repeating trends in physical and chemical properties
Why does first ionisation energy decrease from group2 to group3
in group3, outermost electrons are in p-orbitals whereas in group2 they are in s-orbitals so the electron is more easily removed
why does first ionisation energy decrease from group5 to group6
in group5 the 5 electrons in the p-orbitals are single whereas in group6 the outermost electrons are spin paired
experiences repulsion and more easily removed
does first ionisation energy increase or decrease between the end of one period and the start of another
why?
decrease
increased atomic radius
more shells, more electron shielding
properties of a giant metallic lattice
high melting and boiling point
good conductivity
malleable
ductile
what is a malleable metal
the metal can be shaped into different forms
what is the solubility of giant covalent structures and why
insoluble in almost all solvents
covalent bonds too strong to be broken by interaction with solvents
