bonding and structure Flashcards
definition of ionic bonding?
the electrostatic attraction between oppositely charged ions
what structure do ionic compounds adopt when in a solid state?
a giant lattice structure
why do ions not conduct electricity when in solid state
as ions are not free to move
properties of ionic compounds
- high m.p and b.p
- soluble
- electrical conductors (l/aq)
why do ionic compounds have a high m.p and b.p and what alters it
as ionic bonds are strong - the greater the charge of the ion, the stronger the ionic bond, and therefore more energy required to break it - hence different ionic structures will have different melting/boiling points
why are ionic compounds soluble
the stronger the ionic bond, the less soluble the substance becomes
when are ionic compounds electrical conductors
when they are liquid or aqueous as ions are free to move
covalent bonding def.
the strong electrostatic attraction between a shared pair of electrons and the nuclei of the bonded atoms
which electron doesn’t reach its full outer shell?
boron
dative covalent bonding def.
a covalent bond in which the shared pair of electrons have been supplied by one of the bonding atoms only
result of 4BP + 0LP
tetrahedral
bond angle of tetrahedral shape
109.5
result of 3BP + 1LP
pyramidal
bond angle of pyramidal shape
107
result of 2BP + 2LP
non-linear
bond angle of non-linear shape
104.5
result of 3BP + 0LP
trigonal planar
bond angle of trigonal planar
120
result of 5BP + 0LP (pf5)
trigonal pyramidal
result of 6BP + 0LP (sf6)
octahedral
bond angle of octahedral shape
90
result of 2BP + 0LP (CO2)
linear molecule
bond angle of linear shape
180
electronegativity def.
a measure of the attraction of a bonded atom for the pair of electrons in a covalent bond
most electronegative element
fluorine
least electronegative element
francium
electronegativity trend across a period and down a group
increases as you go across a period, decreases as you go down a group
non-polar bond def.
when the electronegativity difference is 0 or close to 0 eg diatomic molecules / with no charge separation across bond in a molecule
intermolecular forces def.
weak attractions between dipoles on different molecules
list types of IFS
- induced dipole-dipole interactions (London forces)
- permanent dipole-dipole interactions
- hydrogen bonding
when are London forces present
always present in all molecules and atoms
how does boiling point and going down the group affect London forces
- as we go down the group
- boiling point increases
- so no. of electrons increase
- so London forces between molecules are stronger
define a simple covalent lattice in solid state
weak London forces are holding the covalent bonds between molecules
what is a simple covalent molecule
London forces between molecules in gas/liquid state
properties of simple covalent molecules
- low m.p. and b.p.
- low density
- cannot conduct electricity as no free moving electrons
- the polar molecules are soluble in polar solvents, vice versa for non-polar molecules
hydrogen bonding def.
a strong dipole-dipole attraction between an electron-deficient hydrogen atom on one molecule and a lone pair of electrons on a highly electronegative atom containing N, O or F on a different molecule
hydrogen bonding criteria
- electron deficient hydrogen
- lone pair of electrons on N, O, or F
how many hydrogen bonds can one water molecule form?
4
anomalous properties of water
- higher m.p/b.p compared to other simple covalent molecules
- ice is less dense than liquid water
- cohesive eg transpiration in plants
- high surface tension (many h2 bonds)
why is ice less dense than water?
when all the h2 bonds form, they hold the water molecules far apart creating an open lattice structure. when they’re further apart this means they’re less compact, therefore less dense
metallic bonding def.
a strong electrostatic attraction between positive ions in a fixed position and delocalised electrons
what is a giant metallic lattice for
all metals will adopt a giant metallic lattice
properties of a giant metallic lattice
- high m.p/b.p
- conduct electricity due to delocalised/free moving electrons
- insoluble - due to strong metallic bonds
what makes a stronger metallic bond
the higher the charge of the ion
