What is the order of strength of intermolecular, ionic, covalent, network and metallic bonding? and why?
347735
1)ionic and covalent depend on the situation and are considered generally equal \n2)network bonds are a type of covalent bond\n3)metallic bonds are weaker (this one's kinda fuzzy)\n4)intermolecular are the weakest
482712
um...  subtract???  hahaha\n\nThe higher the differences in electronegativity, the more ionic the compound.  The higher electronegative atom/ion will want to "hog" all the electrons.  They are greedy like that.  They don't want to share (which would be a covalent bond).
952034
polar refers to bonds and molecules where the electrons tend to be closer to one atom than another because of the "strength" with which the atoms "pull" the electrons.  This "strength" is usually measured by a number called "electronegativity".  There should be a chart or list of these numbers in your text book.\n\n"Covalent" is a type of chemical bond where the electrons are shared between atoms, rather than transferred (ionic bond).  Usually covalent bonds form between two non-metals.\n\nFrom your question it looks like the first two terms (polar and nonpolar) are referring to individual chemical bonds.  Find the two atoms attached to either end of the bond, look up their electronegativities and see if one atom "pulls" on the electrons stronger than the other.  This would produce a polar BOND.\n\nThe last two (nonpolar covalent; polar covalent) could refer to both a bond or a molecule.  For the molecule you need polar bonds,  and then you have to know the shape of the molecule since it is possible for the individual bond "pulls" to cancel each other - kind of like a tug of war.  Generally if a molecule is symmetrical (triangle, tetrahedron or linear triatomic) the "pulls" cancel.  Although the BONDS may be polar covalent, the entire MOLECULE would be nonpolar covalent.
324098
Hydrogen Bonds, they are the strongest type of intermolecular forces. It occurs between H-F, H-N, and H-O.
59794
i dunno the first one...but sharing electrons is a covalent bond, and charged atoms are ions and they form ionic bonds
31630
Put simply, the term covalent bond is used to describe the bonds in compounds that result from the sharing of one or more pairs of electrons.\n\nCovalent bonding is an intermolecular form of chemical bonding characterized by the sharing of one or more pairs of electrons between two species, producing a mutual attraction that holds the resultant molecule together. Atoms tend to share electrons in such a way that their outer electron shells are filled. Such bonds are always stronger than the intermolecular hydrogen bond and similar in strength to or stronger than the ionic bond.\n\nCovalent bonding most frequently occurs between atoms with similar electronegativities. For this reason, non-metals tend to engage in covalent bonding more readily since metals have access to metallic bonding, where the easily-removed electrons are more free to roam about. For non-metals, liberating an electron is more difficult, so sharing is the only option when confronted with another species of similar electronegativity.\n(Please definitely see the wikipedia site! it has pictures that explain it better)\n\nAlso look at this site for high school level explanations. This page explains what covalent bonding is. It starts with a simple picture of the single covalent bond, and then modifies it slightly:\nhttp://www.chemguide.co.uk/atoms/bonding/covalent.html\n(they have lots of examples)
729609
In covalent bonds, elements actually lose and gain electrons to other elements in the compound.  Polar bonds are covalent bonds that don't require hybridization, and Non-Polar bonds do require hybridization.  Ionic bonds occur when elements share an electron but do not lose it.  Therefore, you can tell that NaCl is ionic because Sodium has an extra electron and Chlorine needs one, so Sodium shares it's electron with Chlorine, forming a very strong bond.  Ionic bonds, however are broken by Polar-Bonded liquids such as water, which is why NaCl dissolves.
188677
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.
233865
Generally, the elements that form ionic bonds tend to have either very high or very low electronegativity. This means the Alkali metals, Alkaline Earth metals, and the Halogens, especially when reacting with each other.
452327
Not exactly sure what you're asking here either, but perhaps a better answer is that electrons that are shared equally between two atoms in a bond are covalent, while electrons that spend more time with one atom than with the other form an ionic bond.\n\nHope that helps.
1121228
I don't see any example of Metallic or network covalent solid materials in the compounds you listed.  \n\nBaCl2 is called Barium Chloride because it is an ionic compound.  You only use Mono, Di, Tri, etc prefixes for molecular compounds.  \n\nLiCl = Ionic\nBaCl2 = Ionic\nBCl3 = molecular\nCCl4 = molecular\nNCl3 = molecular\n\nAlthough Boron can possibly form network covalent bonds much like Carbon and Silicon, the fact that Boron Trichloride has very low melting point as well as boiling point clearly indicates that it does not have network covalent bonds.  A network covalent compound is very stable and has very high M.P and B.P like diamond and quarts. \n\nMetallic bonding is the is the bonding within metals. It involves the localized sharing of free electrons among a lattice of metal atoms.  Which means that to have metallic bonding the compound can not have any nonmetals in it.  In your list of compounds, they all have nonmetals in the compound, therefore non of them are metallic.
291938
1. electron\n2. ionic\n3. molecules of compounds\n4. decomposition (not sure)
496847
N2- has a bond order of 3.\n\nhttp://www.meta-synthesis.com/webbook/39_diatomics/diatomics.html
147054
Metallic bonding is the bonding within metals. It involves the delocalized sharing of free electrons between a lattice of metal atoms. Thus, metallic bonds may be compared to molten salts.\n\ne.g. Iron (Fe)... Why is it so strong? Of course, metallic bonding.\n\nMetal atoms typically contain a small amount of electrons in their valence shell compared to their period or energy level. These become delocalised and form a Sea of Electrons surrounding a giant lattice of positive ions.\n\nMetals seem to have higher boiling and melting points which might suggest stronger bonds between the atoms.\n\nMetallic bonding, as with covalent bonding is non-polar, in that there is no (for pure elemental metals) or very little (for alloys) electronegativity difference among the atoms participating in the bonding interaction, and the electrons involved in that interaction are delocalized across the crystalline structure of the metal.\n\nThe metallic bond accounts for many physical characteristics of metals, such as strength, malleability, ductility, conduction of heat and electricity, and luster. See also chemical bond.\n\nMetallic bonding is the electrostatic attraction between the metal atoms or ions and the delocalised electrons. This is why atoms or layers are allowed to slide past each other, resulting in the characteristic properties of malleability and ductility.\n\nMetallic bond are different from chemical bonds. Ionic and covalent bonds are chemical bonds.\n\nCovalent bonding is a common type of bonding, in which the electronegativity difference between the bonded atoms is small or non-existent. In the latter case, the bond is sometimes referred as purely covalent. See sigma bonds and pi bonds for current LCAO-explanation of non-polar bonds.\ne.g. CH4 (methane)\n\nIonic bonding is type of electrostatic bond between atoms which have an electronegativity difference of over 1.6 (this limit is a convention). These form in a solution between two ions after the excess of the solvent is removed.\ne.g. NaCl (common salt)\n\n<BEST ANSWER>\n============================================================\nHey, I tried to draw the structures of NaCl and methane molecules but this Yahoo! made it impossible. I hope you know the structurel.
1064705
I think knowledge of the terms electronegativity and bond energy address your question. \n\nElectronegativity refers to the difference by which two atoms tend to hold onto electrons. You should be able to find a periodic table that shows you the electronegative values of all the atoms. You can compute the electronegativity of a bond between two atoms by subtracting their values from each other (based on Linus Pauling's calculations). For example, Flourine (F) has an electronegativity value of 4 and Na = .9. The electronegative value of an F-Na bond is 3.1 (4 - .9). My old high school textbook (Heath) states that ionic bonds are those having an electronegative difference between 1.7 and 4; polar covalent bonds are between 0.2 and 1.7; covalent (non-polar) bonds are between 0 and 0.2. Thus an F-Na bond by this classification system is an ionic bond.\n\nBut I suppose to answer your question most directly, I would say the evidence that ionic bonds are strong concern something called bond energy. That is, the difference between how much energy is needed to make and break chemical bonds. Generally speaking, the higher the electronegative difference (which translates into more of an ionic bond--see calculation above), the more energy required to break such bonds. \n\nThe bottom line, is that I think this translates into something like the observed boiling points of various substances. Substances that have strong ionic bonds (again, determined by calculation above) I believe have a higher boiling point (because it takes more energy/heat) to break the bonds.
161637
ionic,the diference in electronegativity values is greater then 1.6.  the metal bonded to a non metal method is also correct
1008726
Strong Love!\n\nIn other words, it's a covalent bond because sharing of electrons takes place in these elements when they combine with each other.