CHEMICAL BONDING AND STRUCTURE Flashcards
Giant structures Discrete molecules Physical properties
metallic bonding
strong electrostatic attraction between the nuclei of metal cations and delocalised electrons.
metals properties
high mp/bp high thermal/electrical conductivity malleable sonorous ductile low IE's
Why do metals have high mp’s?
Need to break strong forces of attraction between the cation nuclei and delocalised electrons.
also giant lattice structure requires much energy to break
What determines bp/mp’s?
no. of delocalised electrons (group 1 is low while d-block metals are higher)
size of cation (smaller radii has closer electrons to nuclei higher mp)
electrical conductivity in metals
movement of delocalised electrons towards the positive terminal of a cell with a potential difference.
thermal conductivity
delocalised electrons passing KE along metal.
malleability
shape configuration
ductility
ability to draw metal into a wire
what does ductility and malleability depend on?
the cation and electron movement ability. layers slide over one another and delocalised electrons prevent strong forces of repulsion between cations.
ionic bonding
strong electrostatic attraction between oppositely charged ions.
arrangement of ionic substances
giant ionic lattice in which electrostatic attraction occurs in all directions.
how do you determine the strength of electrostatic attractions in an ionic substance?
calculation of energy per mole of solid needed to separate ions to infinity in which ions can no longer interact
how is strength of ionic bonding affected by size of cation?
as cation size increases, amount of energy needed to separate them decreases
how is strength of ionic bonding affected by size of anion?
increases as anion size increases
how does size of ions affect energy needed to separate ionic substances to infinity?
smaller substances require more energy to separate
why must data of ionic radii all come from the same source?
because radii is difficult to measure and several methods could conflict in accuracy.
properties of ionic substances
high mp/bp
brittle
water-soluble
poor electrical conductivity as a solid
why do ionic substances have high melting points?
giant lattice networks of oppositely charged ions have combined large electrostatic forces.
why are ionic substances brittle?
as stress causes ionic layers to slide over one another, meaning same charge ions are next to each other, repelling
why are ionic substances water soluble
polarity of water molecues allows separation of ions
reason for electrical conductivity when molten
previously fixed electrons become mobile, allowing a migration of electrons towards the positive terminal
what sort of current is used in electrolysis
direct
covalent bonding
2 atoms overlapping atomic orbitals of which contain a singular electron
3 ways covalent bonds interact
sigma (2 s orbitals)
sigma (2 P orbitals)
pi (2 P orbitals)
how may pi orbitals form? how does this affect the molecule?
sigma bonds may lead to the formation of pi bonds, creating larger electron density above and below the molecule
bond length
the distance between nuclei of 2 atoms covalently bonded.
how does bond length affect covalent bond strength
shorter length, greater strength
electronegativity
the ability of an atom to attract a bonding pair of electrons.
how does electronegativity change over the periodic table?
decreases down groups
increases across periods
What are polar molecule electron densities like?
they form asymmetrical electron densities, resulting in one atom having a slight positive charge and the other a slight negative charge (delta minus and delta plus)
what are polar bonds represented by?
arrows from one atom to the other
2 bonding ideals
100% ionic
100% covalent
polar bond
type of covalent bond between 2 atoms where bonding electrons are unequally distributed, meaning slight differing charges of each atom.
Discrete (simple molecules)
Electrically neutral group of 2 or more atoms held together by chemical bonds.
Dative covalent bond
Bond formed when an empty orbital of one atomoverlaps with an orbital containing a lone pair of electrons.
Dative bond representation
Arrow
Formation of Al2Cl6
Lone pair on chlorine atom overlaps with empty orbital on an aluminium atom.
Only occurs in gaseous state.
VSEPR
Valence shell electron pair repulsion theory
electron pair repulsion theory basis
The shape of an electron/ion is caused by repulsion between electron pairs (lone and bond) that surround the central atom, arranging themselves to experience minimal repulsion.
order of repulsion strength
lone pair-lone pair repulsion > lone pair-bond pair repulsion > bond pair-bond pair repulsion
steps to work out shapes of molecules
1.) find no. of areas of electron density (bond pairs/lone pairs) via dot and cross.
Then work out shape
2 bond pairs and no lone pair shape
angle?
linear
180 degrees
3 bond pairs and no lone pairs shape and angle
trigonal planar
120 degrees
4 bond pairs and no lone pairs shape
angles?
tetrahedral
109.5 degrees
5 bond pairs and no lone pairs shape
angles?
trigonal bipyramidal
90 degrees and 120 degrees
6 bond pairs and no lone pairs shape
angles?
hexagonal
90 degrees and 180 degrees
3 bond pairs and one lone pair
shape? angle?
why?
trigonal pyramidal
107 degrees
lone pair-lone pair repulsion is greater than bond pair-lone pair repulsion so the angles is slightly less than 109.5 degrees.
2 bond pairs and 2 lone pairs
shape?
angles?
v-shaped
104.5 degrees
lone pair-lone pair repulsion > lone pair-bond pair repulsion so bond angles is further depressed.
dipole
a separation of charge caused by the drift of bonded electrons towards the more electronegative element.
why does electronegativity decrease down a group?
as atomic radii increases, distance between electrons and nuclei increases, reducing force of attraction between atoms.
also more shielding
molecule with intermediate bonding character (both covalent and ionic)
calcium carbonate
ionic between calcium and carbonate, covalent within carbonate
rough size of an atom
10 to the -10 m
isoelectronic
having the same electronic configuration
size of negative ions in relation to atoms that they form from
what about positive?
negative ions are larger
positive ions are smaller
what is ionic bonding complicated by?
ionic packing
covalent character
dative bonding
covalent bond in which both electrons come from the same atom
discrete (simple) molecule
electrically neutral group of 2 or more atoms held together by weak chemical bonds
DIPOLE
when 2 charges of equal magnitude but opposite signs are separated by a small distance
factors affecting electronegativity
nuclear charge
atomic radius
shielding
charge effects on electronegativity
higher charge, greater electronegativity
atomic radius effect on electronegativity
smaller atom, closer the pair of bonding atoms are to the nucleus, meaning stronger force of attraction
shielding effect on electronegativity
more inner shells, more shielding and so less force of attraction to attract 2 bonding pairs
electronegativity down a group
decreases
bigger radius
more shielding
less force of attraction
electronegativity across a period
increases smaller radius higher charge less shielding more force of attraction
diatomic molecule polarities
non-polar
as 2 atoms in each molecule are same and so have same eneg w symmetrical densities
symbol representing dipole
arrow w cross through
linear molecules (symmetrical) polarity
non-polar as dipoles cancel each other out
trigonal planar molecules polarity
non-polar as is symmetrical so dipoles cancel each other out
tetrahedral polarity
if symmetrical, non-polar
if unsymmetrical, polar as dipoles reinforce one another
trigonal pyramidal polarity
dipoles reinforce one another so is polar
v-shaped molecule polarity
dipoles reinforce each other so is polar
3 types of intermolecular forces
London/dispersion forces
permanent dipole interactions
hydrogen bonds
formation of London/dispersion forces
created by instantaneous dipoles occurring due to electron density fluctuations, inducing other dipoles between other molecules, aligning so that they interact favourably with one another
why do London forces continue to attract each other despite movement
as fluctuations in electron density are much quicker than the kinetic energy, so continue to attract each other regardless of energy
factors affecting London forces
attractive force increases w increasing number of electrons in the molecule (fluctuation in density faster)
more points of contact between molecules, greater overall London force
permanent dipole interactions
molecules possessing permanent dipoles interact and align, causing forces of attraction between molecules
why is London force more significant than permanent dipole interactions
random movement of molecules mean they have less alignments and so less interactions
sometimes forces repel so contribute little overall
where do london forces exist
between all types of molecules, regardless of polarity
where do hydrogen bonds exist
between all atoms containing an OH group
atom bonded must be more eneg than hydrogen
example of hydrogen bonding
betweenoxygen atom of one molecule and hydrogen of another
h bonding
h bonded to an electronegative atom
reasons for increasing bp with increasing molecular mass
more electrons and so more instantaneous and permanent dipoles form
carbon chain length increase results in more points of contact between adjacent molecules
branched chain effects
lower bp
due to less points of contact so decrease in IMF
anomalous properties of water
relatively high bp for little electrons
density of ice is less than water
2 conditions needing to be met for a substance to dissolve
solute particles must be surrounded by solvent particles
forces of attraction between solvent and solute particles must be strong enough to overcome solute and solute forces
van der waal forces
sum of al the intermolecular interactions between the molecules, including London and permanent dipole-dipole forces.
what happens to the ionic radius of a set of isoelectronic ions as the atomic number increases
decreases as the electrons are attracted to the nucleus more strongly.
giant ionic lattice
repeating units in a regular structure, caused by ions attracted in all directions to oppositely charged ions.
