what typeof bonding holds a polyatomic ion together?
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Covalent bonds, as in NH4+
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A covalent bond\nThe other is an Ionic Bond which is when one atom gives an electron to another atom.
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A chemical bond is the physical phenomenon of chemical species being held together by attraction to each other through sharing, as well as exchanging, of electrons and is a phenomenon that is fully described by the laws of quantum electrodynamics. In general, strong chemical bonds are found in molecules, crystals or in solid metal and they organize the atoms in ordered structures.
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A coordination bond is a bond where one molecule gives both electrons to form a covalent bond. It differs from an ionic bond because the electrons are "shared" not given (e.g. the bond consists of a wavefunction, instead of discrete cations and anions)
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Ionic bond, hydrogen bonds, covalent bond, ionic interactions, van der waals interaction, non-polar interaction ... the list goes on.\n\nA bit more on covalent bonds:\nThere are sigma and pi covalent bonds, mainly. Pi bonds are usually the double bonds, while sigma are single bonds. Pi bonds can have resonance!\n.\nAlso there are often non-bonding electrons and anti-bonding electrons in molecules, but they aren't really "bonds"\n.\nCheers.\n.\nP.S.: If you share "chemistry" with a girl / boy you could also call that a "chemical bond"! He he he he.
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an ion with both acidic and basic part
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1. Crystals\n\n2. Covalent bonds\n\n3.  Molecules\n\n4.  Vasoline\n\nCrossword = nucleuss\n\nH- Ion
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A covalent bond is a bond in which the atoms making up the resulting molecule "share" the electrons in the form of covalent bonds.\nAn ionic bond is an extremely one sided "sharing" of electrons, so one sided, one atom is said to have completely stripped of an electron from another atom.\n\nWhat determines if the bond is covalent or ionic is the difference in the electronegativities of the two atoms bonding.  If the electronegativity difference is high enough, the bond is said to be ionic.  If the difference in electronagativities is significant, but not quite to the level of a full ionic bond, the bond is said to be polar (which may or may not mean that the resulting molecule is polar, this depends on several things).\n\nLooking at a chart of [Pauling] electronegativities,\nhttp://www.psigate.ac.uk/newsite/reference/plambeck/chem1/p01244a.htm\none can see that most elements are somewhere in the 1 to 2 range, with only a few elements getting much lower or higher than this value.  Unless the difference is high enough (usually a difference of about 1.5 is considered ionic), then the bond will just be polar, or less so.\n\nSo, all elements have the potential to theoretically bond ionically, but only in certain combinations with other elements.\nFor example, you will never see an ionic bond between Carbon and Hydrogen since the difference is just .4, but an ionic bond between Lithium and Fluorine is easy to make with a huge difference of 3.0.
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Covalent bond: bond in which one or more pairs of electrons are shared by two atoms.\n\nIonic bond: bond in which one or more electrons from one atom are removed and attached to another atom, resulting in positive and negative ions which attract each other.\n\nOther types of bonds include metallic bonds and hydrogen bonding. The attractive forces between molecules in a liquid can be characterized as van der Waals bonds.
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Alkali metals and alkali earth metals tend to form ionic bonds.  They have a +1 or +2 charge.\n\nThe halogens normally form ionic bonds when "joined" with a IA or IIA metal.\n\nTransition metals normally are ionic.\n\nNow for the trick part:  Carbon and nitrogen generally form covalent bonds.  When joined with other elements, they can produce a neutral molecule (CO, CO2).  But they also form polyatomic ions (C03- (carbonate); NO3- (nitrate) ).  \n\nGeneral rule of thumb - IA, IIA, and VIIA atoms will make ionic bonds.  \n\nHope this helps!
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1) covalent\n2) ionic
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Covalent bonding is a description of chemical bonding that is characterized by the sharing of one or more electrons between two atoms. In general bonds are defined by a mutual attraction that holds the resultant molecule together. Often bonding occurs in such a way that the the outer electron shells of the participating atoms become filled. In contrast to electrostatic interactions labeled as "ionic bonds," the strength of covalent bond depends on the angular relation between atoms in polyatomic atoms. Covalent bonding is most important between atoms with similar electronegativities. Covalent bonding is often delocalized.
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The monomers are linked by covalent chemical bonds.
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Figure 21: Heteropolymolybdate structures. (A) The structure of [Mo< ...  \n  \nchemical compound\n\nCrystal field, ligand field, and molecular orbital theories \n\nConsiderable success in understanding certain coordination compounds also has been achieved by treating them as examples of simple ionic or electrostatic bonding. The German theoretical physicist Walther Kossel's ionic model of 1916 was revitalized and developed by the American physicists Hans Bethe and John H. Van Vleck into the crystal field theory (CFT) of coordination, used by physicists as early as the 1930s but not generally accepted by chemists until the 1950s. This view attributes the bonding in coordination compounds to electrostatic forces between the positively charged metal ions and negatively charged ligands--or, in the case of neutral ligands (e.g., water and ammonia), to charge separations (dipoles) that appear within the molecules. Although this approach meets with considerable success for complexes of metal ions with small electronegative ligands, such as fluoride or chloride ions or water molecules, it breaks down for ligands of low polarity (charge separation), such as carbon monoxide. It also requires modification to explain why the spectral (light-absorption) and magnetic properties of coordinated metal ions generally differ from those of the free ions and why, for a given metal ion, these properties depend on the nature of the ligands. Thus, the complex ion hexafluoroferrate(III) (FeF63-) has magnetic properties to be expected from a substance with five unpaired electrons, as does the free iron(III) ion (Fe3+), whereas the magnetic properties of the closely related hexacyanoferrate(III) ([Fe(CN)6]3-) correspond to only one unpaired electron.\n\nSince 1950 it has been apparent that a more complete theory, which incorporates contributions from both ionic and covalent bonding, is necessary to give an adequate account of the properties of coordination compounds. Such a theory is the so-called ligand field theory (LFT), which has its origin in the more general, but more complicated, theory of chemical bonding called the molecular orbital (MO) theory. (Molecular orbitals describe the spatial distributions of electrons in molecules, just as atomic orbitals describe the distributions in atoms.) This theory accounts with remarkable success for most properties of coordination compounds.\n\nAn important conclusion from LFT is that two types of bonds, called sigma () bonds and pi () bonds, occur in coordination compounds just as they do in ordinary covalent (organic) compounds. The more usual of the two are  bonds, and they are symmetrical about the axis of the bond;  bonds, which are less common, are unsymmetrical with regard to the bond axis. In coordination compounds,  bonding may result from donation of electrons from ligands, such as fluorine or oxygen atoms, to empty d orbitals of the metal atoms (the designation d merely being a way of indicating the particular orbitals involved). An example of this type of bonding is the chromate ion, (CrO4)2-, in which the oxygen atoms donate electrons to the central chromium ion (Cr6+). Alternatively, electrons from d orbitals of the metal atom may be donated to empty orbitals of the ligand. This is the case in the compound tetracarbonylnickel, Ni(CO)4, in which empty  orbitals in the carbon monoxide molecules accept d-orbital electrons from the nickel atom.
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An ion is formed when one electron completely transfers over from one atom to another atom.  The one losing the electron acquires a positive charge and the one losing the electron acquires a negative charge.  But then these two ions are attracted to each other forming an ionic compound.  After coming together, the bonding electrons spend most of the time around the atom with the negative charge, that is, the more electronegative ion.  However, there are also moments when the electrons get closer to the electropositive atom.  Though it's rare to find the bonding electrons near the electropositive atom (since it's orbitals are higher in energy), calculations find that ~5% of the time the electrons are closer to the electropositive atom than to the electronegative atom.  (Don't trust me on the ~5%... this is based on only a few examples from my Inorganic Chemistry chapter I read a few weeks ago).  A popular example is that of HF.  If you use molecular orbital diagrams, you see that the 2s orbital is so low in energy (around -40eV) that it doesn't share it's electrons.  The 2p(z) orbital bonds with the hydrogen's 1s orbital but you'll see that the resulting bonding orbital is so close to the 2p(z) orbital in energy that you can sort of consider it somewhat nonbonding.  The point is that there is little influence from the H 1s orbital, but there is still a small amount of influence.  This means the bonding electrons spend a fraction of the time around the H atom, meaning that the original electron in question has no completely transferred over to the F atom.
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An ionic compound.\nA covalent compound.
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ionic,the diference in electronegativity values is greater then 1.6.  the metal bonded to a non metal method is also correct