Chemistry: What is the difference between a subshell and an orbital?
815589
it is very simple, the shells, subshells actually represent energy levels in the atoms. it is just like big building having different storeys and them again the rooms. You can compare storeys as shells, and individual apartments as subshells and the rooms in each apartment as orbitals where we the electrons are revolving
799741
as i have explained with the formation of ionic compounds. this is an ionic compound like table salt (NaCl)\nLithium gives up the one outermost electron to Chlorine which has a stable octet.\nLithium then has just two. Remember, this is the first orbital (s orbital) with a maximum of two electrons.
806015
three principles that govern this process: \n1) the Aufbau (build-up) principle, \n2) the Pauli exclusion principle, \n3) Hund's rule\n\nExclusion principle \nNo more than one electron can have all four quantum numbers the same. What this translates to in terms of our pictures of orbitals is that each orbital can only hold two electrons, one "spin up" and one "spin down". \n\nBuild-up principle \nYou may consider an atom as being "built up" from a naked nucleus by gradually adding to it one electron after another, until all the electrons it will hold have been added. Much as one fills up a container with liquid from the bottom up, so also are the orbitals of an atom filled from the lowest energy orbitals to the highest energy orbitals. \n\nHowever, the three p orbitals of a given shell all occur at the same energy level. So, how are they filled up? Is one of them filled full with the two electrons it can hold first? Or do each of the three orbitals receive one electron apiece before the any single orbital is double occupied. As it turns out, the latter situation occurs. \n\nHund's rule \nThis rule is applicable only for those elements that have d electrons, and so is less important in organic chemistry (though it is important in organometallic chemistry). It says that filled and half-filled shells tend to have additional stability. In some instances, then, for example, the 4s orbitals will be filled before the 3d orbitals.
923045
True.  Elements are characterized by the nature of their atoms; each atom has a positively charged nucleus and is surrounded by a cloud of electrons.  All atoms of a particular element have the same nuclear charge.  Under normal circumstances, atoms are in a low energy state (the "ground state") in which the shape of the electron orbitals is roughly spherical -- bearing in mind that we're talking about the shape of a fog, without definite edges.  In higher energy states, the electron orbitals can get arbitrarily large.
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Electrons in the outermost shell are called valence electrons.\n·  It is the valence electrons determine an atom’s chemical properties.\n·  Electrons in the inner shells are inner electrons or core electrons.\n\nThe total number of electrons is revealed by the atomic number:\nNa - At#:11 valency:1 valency electrons:1 core:10\nSe - At#:34 valency:2 valency electrons:6 core:28\nPb - At#:82 valency:4 valency electrons:4 core:78
952374
Hey what's up?  Ya I remember learning that about 3 or 4 weeks ago in Chemistry.  S, P, D, F orbitals are the four main sets of orbitals around the nucleus of an atom.  An "S" orbital looks like a sphere, a "P" orbital looks like 2 dumbbells, a "D" orbital looks like a four leaf clover, and a "F" orbital really has no shape, it looks more like a blob, unless viewed by a very powerful microscope.  But anyways, they are also called energy sublevels, which are contained in Principal energy levels.  To better understand what an energy sublevel is, picture this, picture the seats in a wedge-shaped section of a theater.  As you move away from the stage, the rows become higher and contain more seats.  Similarly, the number of energy sublevels in a principal energy level increases as "n" increases.  They are according to the shapes of the atom's orbitals.  Hey one last tip, if you are able to take a look at the periodic table of elements, Group 1A and 2A are "S" orbitals, Groups 3B to 2B are your "D" orbitals, groups 3A to 8A are your "P" orbitals except Helium which is an "S" orbital, and elements 58 to 71 and 90 to 103 are your "F" orbitals.
963808
The ligand will accomodate electrons from the ligands, forming dative bonds.\n\nThe sub-shell also ceases into two groups of orbitals.  One group will be at a higher energy level as compared with the other.  The difference in energy level is acounted by the strength of the ligand, in other words, by its polarising power.  The stronger the ligand, the greater is the energy difference between the two groups of orbitals.
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the answer is a, the number of outer shell electrons it can gain or lose.\n\nthia is because that outer electrons are the ones the define the atom in terms of bonding. \nmass nubmer and atomic number dont matter much [as lower level chemistry] for bonds. we know that the number of molecules it can form is infinite, and inner shells dont matter, leaving us wiht (A), the only right answer
791063
Excited state from what I know is when the electrons violate hund's rule, where the electrons occupy the subshells and fill each one individually. Like in the p orbital there are three subshells each subshell takes two electrons a piece, but they will only fill up completely if all subshells are occupied. This is because of electron repulsion for the like charges(hund's rule). But if the atom fills a subshell completely without first distributing them is when it is said to be excited. I think you'd have to add an electron to boron and have it fill the same subshell in the 2p orbital (either the x,y,or z).
23301
Molecular Orbital theory explains the hybridization of suborbitals in one element, while the Valence Bonding theory describes how electrons are shared or donated in atoms of different elements, to form a compound.
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energy levels/orbitals
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Any atom in main group chemistry can only have a maximum of 8 electrons in its valence orbital (outter shell).\n\nelements want to have this magic number of 8 electrons in there outter shell.  For example the formation of NaCL.  Sodium (Na) has 1 electron in its outter oribital, chlorine (CL) has 7 electrons in its outter shell.  Therefore Na and CL want to form an ionic compund so they both have 8 electrons in there outter shell.\n\nCovalent compounds work in a similar way, except There will be transfer of electrons rather than electron sharing.\n\nThe reason its 8 is because of the S and P orbitals.  there can be a maximum of 2 in the S orbital and 6 in the P orbital.  \n\nThis is an over simplification though as there are orbitals which can hold more than 8 electronsd and f orbitals in the transition metals), but its the basic idea you are taught at your level of chemistry ^^
562475
The electron in 4p orbital is farther than 3p orbital because the electron in 4p belongs to the 4th energy level while the electron in 3p orbital belongs to 3rd energy level. Use Aufbau's principle to answer questions similar to this.
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An electron/energy shell, is a group of atomic orbitals with the same value of the principal quantum number n. \nThus in a silicon atom (1s^2  2s^2 2p^6 3s^2 3p^2)  the number of energy shells with electrons in them is 3.\n\n(just for info: There are 5 occupied subshells (1s, 2s, 2p, 3s, 3p) and 8 occupied orbitals 1s 2s 2px 2py 2pz 3s 3px 3py, out of which the first 6 have 2 electrons and 3px and 3py have one electron each; Hund's rule doesn't allow an electron pair in 3px while the other 3p orbitals are empty)
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Similarities: All S-orbitals are spherical.  Any given molecule has only a single 1s orbital, and a single 2s orbital.\n\nDifferences: S-orbitals grow more complex as you go up to higher levels.  The 1s orbital has no spherical node, while the 2s orbital has 1 spherical node.
261105
Better than what?\n\n\nHere's a description of the result:\n\nEach successive "shell" has the capacity to store one more group of paired electrons in a new kind of orbital.\n\nPrincipal quantum number 1:\n1 s-orbital\n- will hold one pair.\n\nPrincipal quantum number 2: \n1 s-orbital\n- will hold one pair.\n3 p-orbitals\n- will hold 3 pairs.\n\nPrincipal quantum number 3: \n1 s-orbital\n- will hold one pair.\n3 p-orbitals\n- will hold 3 pairs.\n5 d-orbitals\n- can hold 5 pairs, but due to "crowding" wait until the 4th S orbital is filled. Are available to help form double bonds for elements like P, S and Cl though.\n\nFor principal quantum number 4 there is the addition of 7 f-orbitals, which insert themselves after the first d-orbital of the next higher principal quantum number is filled.\n\nAs a result, there are lots of rare metals which form 3+ ions...\n\n\nNow why is this so?\n\nThere is no easy way to explain this with classical mechanic and electrostatic. It has to do with standing waves, elimination by interference, and the wave/matter duality. In short, quantum mechanics. If you don't want to start with that topic, just accept that this is how it works...
887107
B.the electron subshell