Re: Can the orbit of an electron be influenced by another proximal electron?

     To give you a straight answer to your question, yes, the wavefunction 
(and hence the orbital) of an electron can be affected by a nearby 
electron.  You probably know that chemists and physicists use the 
Schrodinger Equation to find the wavefunction for electrons.  However, in 
all but the simplest case, the Schrodinger Equation is unsolvable.  The 
only case (or cases) for which we know the exact solution for the 
Schrodinger Equation occurs for atoms or ions with a single electron (in 
other words hydrogen and any nucleus with one electron).  For all other 
cases, only approximations must be made.

      The reason for the Schrodinger Equation being unsolvable comes from 
the fact that there exists a potential between all the electrons in a 
multi-electron atom.  However, potential between one electron and all the 
other electrons depends on position.  Since we do not know the positions 
of the electrons exactly (only the probability of finding 
them "somewhere") we are forced to make some assumptions and 
approximations.

     The first approximation we can make is rather crude.  It simply 
assumes that there is no potential between any electrons and is called 
the "independent-electron approximation."  We can illustrate this in terms 
of the simplest multi-electron atom, helium.  Since neither electron feels 
the presence of the other, the wave function would just be a spherically 
symmetrical 1s orbital as we see for hydrogen.  (In fact, more accurate 
approximations show that multi-electron atoms exhibit oribtals similar to 
that for hydrogenic atoms - for which we can solve the Schrodinger 
Equation exactly).

     However, with further and more accurate representations, we must take 
the various potential interactions into account, and that brings in a 
concept called shielding.  Shielding simply means that one electron will 
in a sense block the full nuclear charge from another electron.  So 
instead of seeing a full +2 nuclear charge for a helium atom, one electron 
can see anything from 1 < Z < 2 depending on the position of the second 
electron (in this case, Z is the effective nuclear charge, which is the 
nuclear charge felt by an electron due to shielding).  Since the effective 
nuclear charge is somewhat less than the full nuclear charge (the charge 
an electron would see if there were no other electrons present), the 
distribution for the electron we are examining changes.  Specifically, the 
probability of finding the electron father out from the nucleus increases.

     If we want to examine other systems, we can look at ammonia (NH3), 
which definitely shows that electrons do have an effect nearby electrons.  
Since nitrogen is sp3 hybridized in this case, we might predict bond 
angles of 109.5 degrees just as in methane (CH4), where carbon is also sp3 
hybridized.  But the lone pair of electrons present in ammonia repels the 
electrons in the N-H bonds, compressing the bond angle to 107.3 degrees.  
Methane has no lone pairs.  

     The topic of electron-electron interactions is actually a bit complex 
and involves much math.  If you would like, most any college level 
physical chemistry or quantum chemistry book will the basic mathematical 
concepts involved.  I recommend "Atoms and Molecules: An Introduction for 
Students of Physic Chemistry" by Martin Karplus and Richard N. Porter.  
Unlike most books, it shows almost every mathematical step used to reach a 
conclusion in detail.

I hope this helps.