Module 2 - Bonding And Structure Flashcards
define ionic bonding
electrostatic attraction between positive and negative ions
describe how ions are formed in ionic bonding
the transfer of electrons from metal atoms to non metal atoms from the outer electrons ( the ones in the highest occupied principal energy level )
explain why ionic compounds have high melting and boiling points
- ionic bonds are strong and a lot of heat is needed to break them (high energy)
- giant ionic lattice
- strong electrostatic attraction between + and - ions
- larger ionic charges produce stronger ionic bonds, so more heat is required to break the ionic bonds
explain how ionic compounds conduct electricity
- as solids, can’t conduct because their ions are bonded together in the lattice
- when molten, the ions can break free of the lattice and are able to move.
- the ions are charged particles so can carry an electric current
define covalent bonding
the strong electrostatic attraction between a shared pair of electrons and the nuclei of the bonded atoms
explain why simple covalent molecules have low boiling/ melting points
-covalent bonds within the molecules (intra molecular forces) are strong, but the forces of attraction between molecules (intermolecular forces) are weak
- not much energy is needed to overcome these forces of attraction
explain why simple covalent molecules do not conduct electricity
- no ions or free electrons present
explain why giant covalent molecules have high boiling/ melting points
- high energy needed to break because
diamond- 4 strong covalent bonds between every c atom
graphite - 3 strong covalent bonds between every c atom
explain why giant covalent molecules conduct electricity
- delocalised electrons which move to carry charge through the structure
when is a dative covalent bond formed
when one atom contributes both of the electrons needed for the covalent bond
what is needed for a dative covalent bond to form
one atom has to have a lone pair of electrons and the other atom must have a vacant orbital
how is a dative covalent bond represented
as an arrow, which shows the direction of the electron pair donation
define electronegativity
-the ability of an atom to attract the pair of electrons in a covalent bond (electron density)
how is electronegativity measured
on the Pauling scale
what is the most electronegative on the Pauling scale
F (4.0)
what is the least electronegative on the Pauling scale
Fr (0.7)
describe how electronegativity changes on the Pauling scale
-increases right across the period
-increases up the group
what are factors that determine the electronegativity of an atom
-the size of the nuclear charge
-the size of an atom
describe how the size of the nuclear charge affects electronegativity
-the bigger this is, the larger the attraction between the nucleus and the pair of electrons
-the electronegativity goes up
describe how the size of an atom affects electronegativity
-as this increases, the pair of electrons are further from the nucleus, and there will be a shielding effect from the inner electrons
-the electronegativity goes down
explain non polar covalent compounds
-if both of the atoms have equal electronegativity they will both have the same tendency to attract electrons
-the bond is formed roughly halfway between the 2 atoms
explain the bond for atoms with slightly different electronegativity values
-if atom A is slightly more electronegative than atom B, A has greater attraction to the electron pair
-the electron pair will be pulled towards atom a
-the a end of the bond will have more electron density and will become slightly negatively charged
-the b end of the bond will have less electron density and will become slightly positively charged
what does the polarity of molecules depend on
it’s overall shape
what shape makes a molecule non polar
-molecules with overall symmetry
-as the dipoles cancel out because overall the sharing of electrons is equal
describe london forces
-the weakest intermolecular attraction
-form between neighbouring non polar molecules
-only have an affect over a couple of nanometers
describe the formation of london forces
-electrons in an atom constantly move
-at any time one end of the molecule might have low electron density compared to the other
-a temporary dipole has been formed
-the S+ end of the molecule will attract electron density to one end of a neighbouring molecule
-the neighbouring molecule now has a dipole that has been induced by the first molecule - an induced dipole
what happens to london forces as the molecules get bigger
-more surface contact
-more electrons
-stronger london forces
describe permanent dipole dipole forces
-10x stronger than london forces
-happen at the same time as london forces
-form between two polar covalent molecules
-force of attraction between opposite charged on neighbouring molecules
-leads to highs boiling points compared with similar sized compounds that just have london forces
describe the formation of permanent dipole dipole forces
-the highly electronegative atom in a molecule is attracted towards the S+ atom in a neighbouring molecule as the S- tries to attract electron density towards itself
describe hydrogen bonding
-only happens when a H atom is covalently bonded to an N, O or F atom (very electronegative)
-the strongest intermolecular force of attraction as 100x stronger than london forces
-ionic bonds are 100x stronger than H bonds
describe the formation of hydrogen bonds
-a H atom is covalently bonded to an N, O or F atom
-the bond becomes so polarised that the S+ H atom can also form a weak bond with another F, O or N atom in a neighbouring molecule
steps for drawing hydrogen bonds
- always show lone pairs
- always show the polarity
- H bond always from a lone pair
what does the shape of a molecule depend on
-the number of electron pairs (bond pairs and lone pairs) around the central atom in the molecule
describe the properties of a linear molecule
bond pairs - 2
lone pairs - 0
bond angle- 180
what do double bonds count as when working out shapes of molecules
1 pair of bonded electricity
properties of a trigonal planar molecule
bond pairs - 3
can have 2 lone pairs
bond angle - 120
properties of a tetrahedral molecule
bond pairs -4
bond angle- 109.5
properties of a trigonal bipyramidal molecule
bond pairs -5
bond angle - 120 and 90
properties of an octahedral molecule
bond pairs- 6
bond angle - 90
when are irregular shaped molecules formed
-if a molecule or ion has lone pairs on the central atom
-as the shapes are slightly distorted away from the regular shapes
-as there is extra repulsion caused by the lone pairs
properties of a pyramidal molecule
bond pairs-3
lone pairs - 1
total pairs-4
angle-107
properties of a v shaped/ non linear molecule
bond pairs-2
lone pairs-2
total pairs-4
angle-104.5
steps for drawing shapes of ions
- draw outer shell electrons of central atom
- look at charge and either add or remove an electron/ electrons
- pair up the electrons
- work out shape and bond angles
what are the properties of water
-high boiling/ melting point for a simple molecule
-ice less dense than water
why does water have a high boiling/ melting point for a simple molecule
hydrogen bonding between molecules
why is ice less dense than water
-ice has an open lattice with hydrogen bonds holding the water molecules apart
-when the ice melts the rigid hydrogen bonds collapse allowing the water molecules to move closer together
-the distance the molecules are held apart is what makes ice less dense than water
why do molecules have a specific shape with specific angles
-bonds repel eachother equally
-as they contain electrons so will want to be as far apart as possible
how do lone pairs affect the shape and angles of molecules
push bonding pairs closer together
properties of a square planar molecule
bond pairs - 4
lone pairs - 2
90 degrees
why does the bond angle remain unchanged in trigonal planar and square planar molecules when lone pairs are added
lone pairs repel equally from opposite sides
what does a bigger difference in electronegativity result in
-the more ionic a compound is
-a difference of 0 makes it covalent
how do covalent bonds become polar
if the atoms attached to it have a difference in electronegativity
describe a molecule where polar bonds are arranged symmetrically
it is non polar
which element do we put S+ next to
the least electronegative element
which element do we put S- next to
the most electronegative
when does a temporary dipole exist
-when 2 molecules or atoms are near by
-when they move away the dipole interaction is destroyed
how can induced dipole dipole hold some molecules
in crystal structures eg iodine
what happens to induced dipole dipole forces as a molecule gets bigger
-more induced dipole dipole forces
-as there are larger electron clouds
describe induced dipole dipole forces in branched hydrocarbons
-cant pack together as close
-so weakens the induced dipole dipole forces between the chains
-and lowers the boiling point
how do you test for permanent dipole dipole forces in polar molecules such as water
-placing a charged rod near a steady stream of a polar liquid
-liquid should bend towards the rod as the molecules align to face the oppositely charged rod
why are ionic compounds soluble
- usually soluble in water because water molecules have a slight electrical charge so can attract ions away from the lattice
name examples of giant covalent structures
-diamond
-graphite
-silicon dioxide
-silicon
describe the structure of diamond
-tetrahedral arrangement of carbon atoms
-4 covalent bonds per atom
describe the structure of graphite
-planar arrangement of carbon atoms in layers
-3 covalent bonds per atom in each layer
-4th outer electron per atom is delocalised
-delocalised electrons between layers
