bonding and chemical interactions Flashcards
why can elements past period 3 break the octet rule
expand into the d orbital
- However, some of the third-period elements (Si, P, S, and Cl) have been observed to bond to more than four other atoms, and thus need to involve more than the four pairs of electrons available in an s2p6 octet. This is possible because for n=3, the d sublevel exists, and it has five d orbitals.
- Size is also an important consideration: The larger the central atom, the larger the number of electrons which can surround it.
which elements have incomplete octets
H (2), He (2), Lit (2), beryllium (4) and boron (6)
boron
6
beryllium
4
ionic bonding
one or more electrons from an atom with low ionization energy ( energy needed to remove an electron = metal) are transferred to an atom with high electron affinity ( higher= wants more- def = the amount of energy released when atom gains an electron - exothermic) this is usually a nonmetal. Once one steals the others electron(s) - it becomes an anion and is attracted to the cation - this is the bond ( electrostatic attraction)
covalent
shared between 2 atoms
- typically 2 nonmetals that have simular electronegativity
- polarity depends on the difference between the two’s electronegativity
if both electrons are given by one of the atoms in the colvealent bond it is called a
coordinate colvalent
BP of ionic compounds
high - bc of the strength of the electrostatic force
other characteristics of ionic compounds
good conductors, disolve in water and other polar solvents, sold state form a crystalline lattice
bonding force of colvalent bonds
the electrons (both in the shared pair) feel an attraction to the two positive nuclei of the bonded atoms
to remember if nonmetals lose or gain electrons: Nonmetals …..
Nonmetals gain electons to become aNions and Negative
when an ionic bond forms
generally extremely exothermic (energy released)
- even if some energy input is needed (electron affinity- always endothermic) - the amount of energy given off when a salt lattice is HUGE - making in exothermic
properties of covalent compounds
they become there own lil molecules that have weak intermolecular reactions so they are poor conductors, low boiling point sand melting points- don’t break into constituent ions
bond order
number of shared electron pairs between two electrons
- single = 1 bond order
- double = bond order of 2
bond length
shorter with increasing bonds
bond strength
strongest with most bonds ( triple)
dipole moment
p= qd ( q is mag of charge and d is displacement seperating the 2 charges)
lewis base
any compound that will donate 2 electrons to form a coordinate covalent bond
- b’s and d’s ( a Base will Donate 2 electrons)
formal charge
the difference in the valence electrons given in a lewis diagram (when its in a colavent compund) and the actual amount of valence electrons in the neutral atom ( when its not in a covalent compound)
formal charge= V normal - Nonbonding - 1/2 bonded
or
= V normla - dots (nonbonded=2) - sticks (one)
” what’s the difference in the electrons given in the lewis structure and the actual valence electrons that the atom normally has–> if its +1–> this means it has one less electron than what it normally does”
backbone of lewis
H is always on the end
halogens are normally on the end
the next least electroneg is in the center
steps for lewis diagram
- construct the backbone
- tally total valence electrons
- draw single bonds between all atoms present
- complete the octets of the terminal atoms ( remember H has only 2)
- place any extra on the central atom
- and if the central atom is not full octet- add double or triple bonds
resoance structures
its possible to draw coumpounds that have the same atom arrangment bu that differ in the placemnt of the electrons - these are resoanace structures
- actualy distribution is hybrid
formal charges for assessing stability
- if formal charges are present- the one with 2 opposite charges closest together is more stable
- negative charge on themost electronegative atom is best
- smallest number of fc’s
VSEPR theory
geometrical shape (nonbonded pairs want to be as far apart as possible) this is the basis of all the shapes
SF6
octahedral
PCl5
trigonal bipyramidal
NH3
- trigonal pyramidal
- not trigonal planar ( like BH3- which has only 6), bc there is alone paor which repels the 3 other bonds
difference between electronic geometry and molecular geometry
electronic- all electrons (bonded and nonbonded)
molecular- only the bonded pairs ( for this one only needs to consider the cordinate number ( the number of atoms bonded to the central atom)
H2O, CH4 and NH3 electronic geometry
all the same = tetrahedral - bc in electronic geo you treat all electrons ( whether bond or no, they same)
H2O, CH4 and NH3 molecular geometry
molecular geometry depends only on the bonded pairs
- so H2O= bent
- NH3- trigonal pyramidal
- CH4- tetrahedral
the important take away about electronic geometry vs molecular ( bc MCAT mostly concerned with molecular)
the electronic tells us the ideal bond angle - but keep in mind that even though H2O ad CH4 both have the tetrahedral molecular formula–> bc in water there are 2 nonbonded electrons- this creates more push, so the bond angle is less than the ideal
atoms involved in H-bonds
pick up the FON
london dispersion forces get bigger when
the size of the atom increases
