Bonding and Intermolecular Forces Flashcards
lewis dot structures
each dot represents the valence electrons, we place dots around the symbol of the element . one on each side and if there are more then 4 , then we start to pair them up.
electrons in the d subshells
these are not considered valence electrons for the transition metals since valence electrons are in the highest n level
single bond
2 electrons between atoms
lone pairs:
non bonding electrons (unshared pair of valence electrons)
double bonds and triple bonds use how many electrons ?
double bonds use 4 electrons
triple bonds use 6 electrons
what is the formal charge?
helpful way to evaluate a proposed lewis structure is to calculate the formal charge of each atom in the structure. the formal charges won’t give us the actual charges but they tell us if the atoms are sharing their valence electrons in the best way possible, which will happen when the formal charges are all zero.
what is the formula for calculating the formal charge?
FC= V-1/2 B-L
= Valence electrons - 1/2 bonding electrons - lone pairs.
the best lewis structures have?
octet of electrons and a formal charge of zero on all the atoms. if there is a charge, the best structures have negative formal charges on the more electronegative element.
resonance
accurately depict bonding in a molecule.
these structures are depicted when there are double or triple bonds in a molecule with one or more lone Paris of electrons.
what is bond dissociation energy?
energy required to break a bond homolytically.
in homolytic bond cleavage, one electron of the bond being broken goes to each fragment of the molecule. – so two radicals form.
heterolytic bond cleavage:
both electrons of the electron pair that make up the bond end up on the same atom : forms a cation and an anion
these processes are diff and therefore have diff energies associated with them.
for similar bonds, the higher the bond order , the shorter and stronger the bond.
a single bond has a bond order of 1
while a triple bond will have a bond order of 3
bond length/ bond dissociation energy comparisons should only be made for similar bonds. so compare C-C bonds with other C-C and c-o with other c-o bonds.
bc of varying radii.
when comparing the same types of bonds, the greater the s character
the shorter the bond
also , the longer the bond, the weaker it is . the shorter the bond, the stronger it is.
covalent bonds
each atom contributes one or more of its unpaired valence electrons.
the electrons are shared by atoms to help complete both octets.
polarity of covalent bonds
a bond is polar if the electron density between the two nuclei is uneven. this occurs if there is a difference in electronegativity of the bonding atoms, and the greater the difference the more uneven the electron density and the greater the dipole moment.
non polar bonds
electron density between the two nuclei is even. little or no diff in electronegativity.
coordinate covalent bonds
when one atom will donate both of the shared electrons in a bond. ex would be : F3BNH3
here N donates both electrons , therefore its the Lewis base.
F3B will be the Lewis acid as its accepting the electrons
ionic bonds
gaining or loosing electrons.
electrostatic interactions between the cation and the anion.
the strength of the bond is proportional to the charges
on the ions and the strength will decrease as the ions get farther apart or if the ionic radii increase.
VSEPR theory
valence shell electron pair repulsion theory
electrons repel one another, electron pairs weather they are bonding or non bonding will attempt to move as far apart as possible.
therefore, the Total number of electron GROUPS on the central atom of a molecule determines its bond angles and orbital geometry.
what are electron groups?
any type of bond (single, double, triple) and lone pairs of electrons
double and triple count ONLY as one electron group.
the shape of the molecule is called the molecular geometry.
however, when lone Paris are present on the central atom of molecule, the shape is not the same as the orbital geometry.
what is the orbital geometry with ZERO lone pairs on the central atom?
linear if 2 electron groups
if 3 electron groups: trigonal planar
if 4 : tetrahedral
what if there is 1 lone pair of electrons?
1 lone pair and two other electron groups: bent
1 lone pair and three other electron groups: trigonal pyramid
what if there are 2 lone pairs?
if two lone Pairs and two other electron groups: bent
s and p orbital combine
you get sp hybrid orbitals ( 2 hybrids as 2 orbitals originally combined).
determining the hybridization for most atoms
add the number of the attatched atoms to the number of the non bonding electron pairs. ( #atoms + # lone pairs )= electron groups.
if the electron groups are 2 :
sp hybridization
bond angle is 180
orbital geometry is linear
if the electron group is 3
hybridization is sp2
bond angle is 120
trigonal planar
if the electron group is 4
sp3
109.5 degrees
tetrahedral
sigma bonds
bond consist of two electrons that are localized between two nuclei
end to end overlap whereas pi bonds are side to side overlap
pi bonds
is formed by the proper , parallel, side to side alignment of two unhybridized p orbitals on adjacent atoms.
in any multiple bonds there is one sigma bond and the remainder are pi bonds
single bond : 1 sigma bond
double bond : 1 sigma bond and 1 pi bond
triple bond : 1 sigma bond and 2 pi bond.
what are intermolecular forces?
they are relatively weak interactions that take place between neutral molecules.
polar molecules are attracted to ions producing
ion-dipole forces
what are dipole to dipole forces?
attractions between the positive end of one polar molecule and the negative end of another polar molecule.
what is the strongest dipole-dipole force?
hydrogen bonding
what is momentary dipole induced dipole force?
a permanent dipole in one molecule may induce a dipole in a neighbouring non polar molecule .
London dispersion forces
instantaneous dipole in non polar molecule may induce a dipole in a neighbouring non polar molecule.
these are VERY weak transient interactions between the instantaneous dipoles in non polar molecules.
every atom with electrons will experience this ! as the size of the molecule increases the # of electrons will also increase therefore the dispersion forces will also increase
despite being weak, all intermolecular forces including London dispersion forces can have PROFOUND impact on the physical properties of a molecule.
substances with stronger intermolecular forces will experience a greater melting point, greater boiling point, greater viscosity and lower vapour pressure.
dipole forces, hydrogen bonding and London forces are all collectively known as
van der Waals forces!!!
hydrogen bond
in order for a hydrogen bond to form, two criteria have to be met:
1) a molecule must have a covalent bond between H and either N,O, F
2) another molecule must have a lone pair of electrons on an N,O,F atom.
vapor pressure
is the pressure exerted by the gaseous phase of a liquid that evaporated from the exposed surface of the liquid.
the weaker the substances intermolecular forces , the higher its vapour pressure and more easily it will evaporate.
while vapour pressure is determined by the intermolecular forces , what else will influence it?
temperature ! : it will inc with the temperature
if you inc the kinetic energy of the particles (which is proportional to temp) allows them to overcome the intermolecular forces and inc the particles that can move to the gas phase.
ionic solid
held together by electrostatic interactions between cations and anions in a lattice structure.
ionic bonds are strong
these will be solid at room temp ! ex NaCl.
netword solids
atoms are connected in a lattice of covalent bonds
these are very strong and they are VERY HARD SOLIDS at room temp.
ex) DIAMOND
metallic solids
these have freely roaming valence electrons : which are called conduction electrons.
due to this, metals conduct electricity and heat excellently and are malleable and ductile!
almost all metals are solid @ room temp
molecular solids
held together by 3 types of intermolecular interactions
hydrogen bonds
dipole-dipole forces
London dispersion forces
these forces are significantly weaker then ionic, network, or metallic bonds
molecular compounds have much lower melting and boiling point then the other types of solids.
these are often liquids or gas @ room temp
