Week Five Flashcards
How electrons arrange
arrange themselves as far away as possible
Lewis model - deficiencies
H2 and F2 are treated equivalently by lewis dot model
Valence bond theory
assumes electrons in a molecule occupy atomic orbitals
presumes electron density is equally shared
sigma bond
a single bond that is formed by the overlap of the 2 1s shells
overlap of a s and p orbital
sigma bond
every first bond
sigma bond
Hybridisation
when covalent bonds are formed atomic orbitals mix or hybridise to form new sets of orbitals
how many orbitals in hybridisation
must have the same amount as started with
pi bonds
formed by the sideways overlap of 2p orbitals
electron density - pi bonds
above and below the axis joining the two atoms
sigma bond strength
stronger than pi bonds because sideways overlap is not as effective as the end on bond
single bond
sigma
double bond
one sigma one pi
triple bond
one sigma and 2 pi
sp
linear
sp2
trigonal
sp3
tetrahedral
sp3d
trigonal byprimidal
sp3d2
octahedral
Molecular orbital theory - predictions
accurate structures
dissociation energies
line positions
paramagnetism and diamagnetism
molecular orbital theory - molecular bonds
treat them as a sharing of electrons between 2 nuclei
treats as delocalised, spread out
bonding orbital
has lower energy so more stable
fills first
anti bonding orbital
has higher energy less stable
will fill last
sigma orbital
bonding molecular orbital with cylindral symmetry about an inter nuclear axis
bond order
stability of molecule
1/2 ( electrons in bonding orbital - number in anti bonding orbitals)
net bond order of 0
wont form bond
diamagnetic
no unpaired spins, cant interact in a magnetic field
paramagnetic
has unpaired spins
Heteronuclear diatomics
if they have same electronegativity they have similar diagram to homonuclear
Acids
sour taste
corrosive
conduct electricity
bases
bitter
soapy
corrosive
conduct electricity
strong
dissociates fully in water
weak
partially dissociates
Arrenhius acids and bases
form solvated hydrogen cations written as H+ or H3O+
Arrenhius bases form hydroxide OH-
Bronsted Lowry acids bases
acids are protons donors
bases proton acceptors
Lewis acid and bases
acids act as electron pair acceptors
bases act as electron pair donors
Monoprotic acids
acids ionise to form one H+ ion for every acid molecule
Diprotic acids
produces 2H+ ions for every acid molecule
Kw
equilibrium constant
[H3O+][OH-]
ph
-log[substance]
strong bases
fully dissociate in water to give metal cation and hydroxide ions `
neutralisation reactions
when acids and bases are mixed
pOH
-log[OH-]
pH + pOH
14.00 at 25 degrees
acidity constant
[HA]
Acid constant equation
HA +H2O => H3O+ + A-
Basicity constant
[B]
Basicity constant equation
B + H2O => BH+ + OH-
Ka + Kb
Kw
tha larger the Ka value
stronger the acid
the smaller the pKa value
stronger the acid
