3.1.3 Bonding Flashcards
Ionic bonding definition
he electrostatic forces of attraction between oppositely charged ions
Properties of ionically bonded compounds
Solid at room temp
Giant structures
High melting points - the energy must break the ionic bonds in the lattice
Conduct electricity when molten or dissolved un water as they are free to carry current
Brittle
Covalent bonding definition
Electrostatic forces of attraction between the shared electrons and the positive nuclei of the other atom
Ionic bonding affected by
AR
Nuclear charges
Co ordinate / dative bond
One atom provides both electrons in the covalent bond
Have exactly the same strength and length as ordinary covalent bonds between same pair of atoms
Metallic bonding definition
Outer main levels of the atom merge causing delocalisation
The electrostatic forces of attraction between the positive metal ions i and the sea of delocalised electrons
Properties of metals
Metals are good conductors of heat and electricity
Good conductors of heat - as energy spread by increasingly vigorous vibrations of the closely packed ions
Strong
Malleable
Ductile - still in the exact same environment
High melting point
What is strength of metals determined by
Charge of the ion - the greater the charge of the ion the greater the number of delocalised electrons = stronger the electrostatic forces of attraction
Size of the ion - smaller the ion the closer the electrons are to the positive nucleus and the stronger the bond
The number of delocalised electrons - stronger the bond
Melting point of a metal affected by
Amount of delocalised electrons
The charge of the nucleus
ionic compounds structure - effected by
Effected by the size or charge of the ion
Types of lattice formed - depends on the size fo the pos and neg ions which are arranged in alternate fashion
Linear
2bp
0lp
180 degrees
Trigonal planar
3bp
0lp
120 degrees
Tetrahedral
4bp
0lp
109.5 degrees
Trigonal pyramidal
3bp
1lp
107 degrees
Bent
2bp
2lp
104.5 degrees
Trigonal bipyramidal
5bp
0lp
120 and 90 degrees
Octahedral
6bp
0lp
90 degrees
How to explain shape
State number of bonding pair and lone pair electrons
State that electron pairs repel and try to get as far as possible
No lp then state that electron pairs repel equally - lp repel more than bonding pairs
State the shape and bond angle
Lp effect on bond angle
Reduce the bond angle by 2.5
Electronegativity
The power of an atom to attract the bonding pair of electrons in a covalent bond towards itself
Electron density
Describes the distribution of neg charge in a molecules
What does electronegativity depend on
Nuclear charge
Distance between the nucleus and the outer shell electrons
The shielding of the nuclear charge by electrons in inner shells
Trends in electronegativity
Down the group = decreases - more shielding
Across a period = increases - nuclear charge increases - smaller atom
Polarity of covalent bonds
Unequal sharing of electrons between atoms that are covalently bonded together
Affected by electronegativity - more electronegative attract the electrons towards itself in a double bond
Greater difference in electronegativity = more polar
really polar= some level of ionic character
Small difference = purely covalent
PD - PD forces
The attractive forces between two neighboring molecules with a permanent dipole
Polar molecules have permanent dipole
Have a pos and neg charged end
ID - ID forces
Also called van der waals
Not affected by electronegativity
Electron charge clouds are constantly moving
Causes a temporary dipole which induced a dipole on neighbouring molecules
Means the
Hydrogen bonding
Type of PD PD forces
Takes place in NOF molecules - as have a lage enough difference in electronegativity when bonded to hydrogen - bond becomes polarised
H becomes delta neg meaning it can bond with the lp of electrons in neighbouring NOF
Water
Hydrogen bonding in water causes anomalies
High mp and bp
Lots of energy required to break the hydrogen bonds
Id - id - affected by electrons
High surface tension
The ability of liquid surface to resist any external forces
External molecules form hydrogen bonds with the internal molecules
Meaning molecules pull downwards the surface molecules causing the surface of them to become compressed and more tightly together at the surface
Increases tension
Density
Solids more dense than liquids - solid is more closely packed together
Solids are in an open lattice
Bond length remains constant in solid
Tetrahedral shape
key phase when describing shapes with no lone pairs
equal repulsion force
macromolecular
giant molecule with covalent bonding
intermolecular force between one ammonia and one water
hydrogen bonding
why a substance is almost insoluble in water
no hydrogen bonding with the water
PH3 molecule
looks like ammonia
bond between the PH3 molecule and H+
dative/ cordinate bond
pair of elctrons in the PH3 are donated to the H+
squre planar molecule
XeF4 - as the dipoles calncel out - non polar
(Cu(Oh)2)2+
what statement about inorganic ionic compounds is always correct
they form giant ionic structures
how to workout which molecules have all the atoms in the same plane
present of a pi bond means no rotation
meaning alkenes have to have all atoms in the same plane
ammonia fact
has permanent dipoles
Kevlar bonding between chains
hydrogen bonding
when drawing the shape of an ion key point
MUST OUT SQUARE BRACKETS AND THE CAHRGE
ionic bonds strength
ionic bonds strongest
stronger than metallic bonding
ketones an aldehydes and hydrogen bonding
do not have hydrogen bonding !
do have dipole dipole
working out the shap of a charged compound
- workout the numb of valence electrons
- workout what the charge dies to the number of valence electrons
- workout the number of bonding pair
- workout how many left - lpss
square planer
4bp, 2lp
same element attached
repelling equally
bond angle of 90
