bonding and structure: content Flashcards
what is an ionic bond in terms of electronegativity
when the difference between electronegativities is greater than 2.1 and the electron is transferred from the [metal] to the [nonmetal]
structure of ionic substances
ALWAYS giant ionic
strong
physical properties of ionic compuonds
high melting and boiling points
do not conduct in the solid state (no mobile charge carriers)
conduct when aqueous or molten (mobile ions)
most are soluble (polar substances dissolve in polar substances [water])
what factors influence the strength of ionic bonding
the greater the ionic charge
the smaller the ionic radius
the greater the difference in electronegativity
octet rule
atoms want a noble gas electron configuration (eight valence electrons) in order to become stable
**exceptions to the octet rule! **
molecules where the central atom is stable with <8 outer shell electrons
BF₃: boron trifluoride
- boron has six valence electrons
AlCl₃: aluminium chloride
- aluminium has six valence electrons
when can you expand the octet
when the central atom is in group three or higher (max no. of e⁻ in a shell = 2n²)
**exceptions to the octet rule! **
molecules where the central atom is stable with >8 outer shell electrons
PCl₅: phosphorus pentachloride
- phosphorus has 10 valence electrons
SF₆: sulphur hexafluoride
- sulphur has 12 valence electrons
criteria for a dative covalent bond
the atom donating must have a lone pair
the atom receiving must have a vacant orbital
how are dative covalent bonds represented
donor species → receiving species
order or electron pair repulsion
LEAST REPULSION
bonding region - bonding region
lone pair - bonding region
lone pair - lone pair
MOST REPULSION
VSEPR structure: bond going into the page
VSEPR structure: bond coming out of the page
2 bonding regions
0 lone pairs
linear 180°
2 bonding regions
1 lone pair
nonlinear 117.5
3 bonding regions
0 lone pairs
trigonal planar 120°
4 bonding regions
0 lone pairs
tetrahedral 109.5°
3 bonding regions
1 lone pair
trigonal pyramid/pyramidal 107°
2 bonding regions
2 lone pairs
nonlinear 104.5°
6 bonding regions
0 lone pairs
octahedral 90°
4 bonding regions
2 lone pairs
square planar 90°
effect of each lone pair present on the bond angle
reduces the bond angle by roughly 2.5°
methods for identifying molecular shapes/bond angles
- dot and cross diagram
- count e⁻ pairs
- any remaining e⁻ (from group no. of central atom) will be lone pairs
- group no. of central atom
- add one for each bond made
- divide by two for the number of e⁻ pairs
effect of double/triple bonds on bond angles
none: all bonding regions have equal repulsion regardless of bond strength
electron structure of diatoms
nuclei are identical so attraction to electrons is equal
** electron density is distributed equally** : nonpolar
both atoms have the same electronegativity
factors affecting electronegativity
- atomic radius of the atom
- nuclear charge (no. of protons)
- number of principal energy levels
how does atomic radius impact electronegativity
smaller radius → electrons can get closer to nucleus → greater attractive force → more electronegative
how does nuclear charge impact electronegativity
more protons → stronger attractive forces between nucleus and attracted e- pair → more electronegative
how does the number of principal energy levels impact electronegativity
more energy shells → more shielding → weaker attractive forces → less electronegative
most electronegative atoms
FLUORINE, oxygen, nitrogen, chlorine
electron distribution within a polar bond
electronegativities are not equal
electron attraction is not equal
electron density distribution is not equal
one side is slightly more positive than the other: opposite sides create a dipole
polarity of C-H bonds
nonpolar
difference between polar molecules and polar bonds
polar molecules have overall [partial] charges
polar bonds have dipoles but these may cancel out within a nonpolar molecule
when will polar bonds form part of a nonpolar molecule
when the molecule is symmetrical
equal pull on the central atom in [horizontal and vertical] components
all non-central atoms have equal polarity and electronegativity
what holds simple molecules together
intermolecular forces (NOT covalent bonds)
london / induced dipole-dipole forces
experienced by all molecules
weakest IMF
how do London forces work
- electrons move randomly within the e⁻ cloud
- this creates temporary dipoles
- these induce dipoles in neighbouring molecules (like charge repulsion).
- the opposite dipoles attract
what factors affect the strength of induced dipole-dipole forces
Mᵣ of the molecule (bigger → more electrons)
surface area contact between molecules
how does molecule size affect the melting point and why
the larger the molecule, the stronger the IMFs so the more energy is needed to overcome them; hence the higher the melting/boiling points
what imfs do nonpolar molecules have
ONLY idd/London forces
why are permanent dipole-dipole forces stronger than London?
the dipoles are always there
when do you get permanent dipole-dipole forces
between polar molecules
(attraction between ẟ⁺ and ẟ⁻) on adjacent molecules
which atoms can form hydrogen bonds
F-H
O-H
N-H
how do hydrogen bonds work
F,N,O are the most electronegative atoms
there is a big difference in electronegativities so they attract the electron pair very strongly
the hydrogen nucleus is left very ẟ⁺
a lone pair on F, N or O on a neighbouring molecule is attracted to the hydrogen
features of a hydrogen bond (MUST show + label)
partial charges on all atoms
lone pairs clearly displayed
hydrogen bond between lone pair and ẟ⁺ H
why is ice less dense than liquid water
h-bonds
molecules are held further apart in the solid state than in the liquid
fewer molecules per unit volume
more space between molecules (‘open lattice’)
why does water have a relatively high melting and boiling point
anomalously high (when looking at group 6 hydrides) due to the presence of hydrogen bonds
