Bonding And Structure Flashcards
ionic bonding
electrostatic forces of attraction between oppositely charged ions
factors effecting strength of ionic bonds
- ion charges
- ionic radii
- effects the strength of attraction between oppositely charged ions
formation of ions
loss or gain of electrons
isoelectronic ion radii effects
Electron configuration remains constant
- change in nuclear charge
- change in strength of attraction between nucleus and electrons
isoelectronic
ions that have the same number of electrons
trend in atomic radii down a group
- radii increases
- ions have the same charge
- atoms get larger as they gain shells
physical properties of ions as evidence for ions
- high melting and boiling points indicate the strong electrostatic forces existing between the oppositely charged ions in ionic compounds
- ability to conduct electricity when molten or ion solution shows that mobile charge carrying particles exist in these substances
dashed line in a displayed formula
going into the page
blocked out triangle in displayed formula
coming out of the page
migration of ions as evidence for ions
a rectangle of filter paper soaked in water is placed on a slide, electrodes connected and a crystal of ionic material placed between the electrodes
covalent bonding
electrostatic force of attraction between nuclei and shared pair of electrons
dative covalent bonding
when both electrons that are shared come from one atom
physical properties of simple covalent
weak intermolecular forces between the molecules
Little energy is needed to overcome these forces
melting and boiling points are low
physical properties of giant molecular
strong covalent bonds
require a large amount of energy to overcome
high melting and boiling point
conductivity of covalent molecules
do not have free electrons
non-conductors of electricity
Graphite is exceptional in that it does have free electrons and so is a conductor
Single covalent bond strength
- change in atomic radii
- change in distance between nuclei and shared electron pair
- change in the strength of a attraction between nuclei and shared pair of electrons
sigma bond
exists directly between the two atoms; it lies symmetrically along an axis joining the two nuclei
pi bond
made up of two halves forming a double bond
expanding the octet
the process of splitting electron pairs means that an atom in a compound can have more than 8 electrons
electron pair repulsion theory
electrons are all negatively charged and so will repel each other each electron region takes up a position to minimise repulsion
electronegativity
ability of an atom with a covalent bond to attract the bonding pair of electrons
polarisation
one atom will have a greater electronegativity than the other
will have a greater pull on the electrons
distorting the electron region
cause of polar bonds
one atom has a higher electronegativity
one atom become more negative
electrons are closer to that atom
three types of intermolecular forces
London forces
permanent dipole
hydrogen bonds
testing polar molecules
place a electrostatically charged rod next to jet of liquid and see if the jet attracts or repels
dipole
The separation of charge which exists in a polar molecule
permanent dipole
In a polar material there is an attraction between the positive charge in one molecule and the negative charge in the other
London forces
- instantaneous dipole in one molecule
- induces a dipole in the other molecule
- the delta + of one attracts the delta - of the other
- more electrons per atom/molecule yields stronger London forces
strength of london forces
more electrons per atom/molecule
order of strength of IMF
weakest=london forces
permanent dipole
strongest=hydrogen bonds
which elements does H-bonding occur in and why
Nitrogen, fluorine and oxygen
high electronegativity
lone pair of electrons
region of negative charge is attracted towards the exposed proton on the hydrogen in another atom
where are each IMF found
LF in all substances
permanet dipole in polar substances
H-bond in compounds of NOF with H
trend of boiling points in alkanes
increase with molecular mass
more electrons so stronger london forces
branched alkanes have lower boiling points as less contact points for bonds
trend of boiling points of alcohols
The –OH group in alcohols causes hydrogen bonding between the molecules. Therefore an alcohol will have a much higher boiling point than an alkane with a similar number of electrons
trend of boiling point of hydrogen halides
increase down the group with molecular mass
LF stronger
HF forms h-bonds so highest BP
high bp means high or low volatility
low volatility
physical properties of water
high mp and bp
can form two hydrogen bonds
what happens when water freezes
hydrogen bonds form an open lattice
low density
ice is on the surface acting as an insulator
water beneath remains liquid
solubility and IMF
if solvents and substances have similar IMF they tend to dissolve
solubility of ionic compounds in water
polar water attracts ions=hydration
water is shaking so as they bond the ions break free from the lattice
sometimes electrostatic forces are to strong to break
solubility of organic compounds in water
Alcohols are soluble can form hydrogen bonds with water
water will not dissolve substances which are not ionic or able to form hydrogen bonds
the dissolving process involves molecules of the dissolving substance to intersperse themselves between the water
strong hydrogen bonds prevent other molecules from moving between water unless they are able to form equally strong interactions with water
solubility in non-polar solvents
dissolve non-polar substances
as have similar IMF so can interact
metallic bonding
electrostatic force of attraction between positive metal ions and delocalised electrons
physical properties of metallic bonding
delocalised electrons allow for conduction of electricity and heat
absence of fixed bonds allows the ions to move layers slide over each other=malleability and ductility
factors affecting strength of covalent bonds
number of electrons shared
distance between nuclei
factors affecting the strength of metallic bonds
- change in ionic charges
- change in ionic radii
- delocalised electrons per ion
- change in strength of attraction between cations and delocalised electrons
type of structure in ionic
giant lattice
strong electrostatic attraction
types of structure in covalent
giant lattice strong covalent bonds
simple molecular strong covalent bonds with weak IMF
type of structure in metallic
giant lattice
strong electrostatic attraction
structures formed by carbon
graphite
diamond
graphene
physical properties of diamond
strong covalent bonds
high mp bp
non-conductor
hardest natural substance
physical properties of graphite
strong covalent bonds and LF high mp bp one electron free per atom as only three bonds=conductor layers can slide=lubricant non-reactive=electrolysis electrode
Multiple covalent bond strength
- change in number of shared electron pairs
- change in distance between nuclei and shared electron pairs
- change in strength of attraction between nuclei and shared electron pairs
Explain shape with all bonds
- minimise repulsion between electron pairs
- number of bonding electron pairs
Explain shape with lone pairs
- minimise repulsion between electron pairs
- number of bonding and lone electron pairs
- lone electron pairs repel more than bonding electron pairs
Electronegativity across a period
- distance and shielding between nucleus and electron pair is constant
- change in nuclear charge
- change in attraction between nucleus and electron pair
Electronegativity down a group
- change in distance and shielding between nucleus and electron pair
- change in nuclear charge
- change in the attraction between nuclei and electron pair
Hydrogen bonds
O/N/F are very electronegative
- causing the bond with H very polar
- H delta + attracts lone pair of ONF from other molecule
