could you explain me about covalent bands ?
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Force holding atoms in a molecule together as a specific, separate entity (as opposed to, e.g., colloidal aggregates; see bonding).\n\nIn covalent bonds, two atoms share one or more pairs of valence electrons to give each atom the stability found in a noble gas. In single bonds (e.g., H-H in molecular hydrogen), one electron pair is shared; in double bonds (e.g., O=O in molecular oxygen or H2C=CH2 in ethylene), two; in triple bonds (e.g., HCºCH in acetylene), three. In coordinate covalent bonds, additional electron pairs are shared with another atom, usually forming a functional group, such as sulfate (SO4) or phosphate (PO4). The number of bonds and the atoms participating in each (including any additional paired electrons) give molecules their configuration; the slight negative and positive charges at the opposite ends of a covalent bond are the reason most molecules have some polarity (see electrophile; nucleophile). Carbon in organic compounds can have as many as four single bonds, each pointing to one vertex of a tetrahedron; as a result, certain molecules exist in mirror-image forms (see optical activity). Double bonds are rigid, leading to the possibility of geometric isomers (see isomerism). Some types of bonds, such as the amide linkages that join the amino acids in peptides and proteins (peptide bonds), are apparently single but have some double-bond characteristics because of the electronic structure of the participating atoms. The configurations of enzymes and their substrates, determined by their covalent bonds (particularly the peptide bonds) and hydrogen bonds, are crucial to the reactions they participate in, which are fundamental to all life
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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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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.
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Try doing a Google search under the Hall or the Hall-Heroult process.    That or check out a General or an Inorganic Chemistry book from the library.  That's a good launch point.
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This is kinda hard to draw on here...\n......_\nK--Br|\n......_\n\nThe lines around Br represent 2 electrons, and the -- represents a bond...ignore the periods
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In covalent, atoms share electrons, but in ionic, one atom gives them to the other.  So in an ionic compound, there is a positive ion and a negative ion.  They ions can separate when put in water, and then the electricity can pass through.  Covalent compounds don't have a positive pole and a negative pole, and they don't ionize in water.  Also, the ionic bond is kinda stronger, so usually the compound is solid at room temperature (salt is ionic, oil is covalent).  You have to know that there is a third category, a polar molecule (like water).
11021
covalent bonds are called the directional bonds.because the orbitals involved in bonding have specific direction where they bond.
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what you need to do, get a SI data book, and look up their natural state. not bonded with anything else, and possibly itself.\n\nPhophorus - white: molecular (P4). red: Chain. Black: Layer\nIodine - molecular (I2)\nMagnesium - hexagonal closest-packing\nNeon - cubic closest-packing\nArsenic -grey: 3d structure - Layer formation. yellow: molecular As4\nSulfur- molecular S8\nBoron - complicated structure.\nSelenium - grey: chain, Red: molecular: Se8\nGallium - Complicated\nCesium - body-centred cubic\nOxxygen - O2, O3 molecular
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See if this article has a diagram\n\nhttp://jchemed.chem.wisc.edu/Journal/Issues/1999/Apr/abs542.html
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Basically CO is a very good ligand to transtional metals because of it great pi accepting properties of it anti bonding orbital with it interation with metal d orbitals. \n\nTherefore any transitional metal form good absorber with CO, example is that of Fe.
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covalent molecules\n-share electrons with each other\n-weak Van der waals forces between molecules\n-strong electrostatic force between atoms\n-low melting and boiling point\n\ndiatomic molecules (group 17)\n-acidic\n-can react with water to form two acidic solution\n-can react with alkali metal to form halide salt\n-reactivity decreases
59794
i dunno the first one...but sharing electrons is a covalent bond, and charged atoms are ions and they form ionic bonds
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I'm sorry...my answer to this question, will either sound really smart or really stupid. \n\nMy AP Chemistry class just finished the chapter on bonding. We learned of Ionic bonds, Covalent Bonds, and Hydrogen bonds. \n\nAn ionic bond is a bond between a metal and non-metal and involves the transferring of electrons so that both atoms will become stable but with opposite charges, become attracted to each other, and form a molecule. An example is NaCl or table salt. Na is a metal, Cl is a non-metal.\n\nThe next is covalent bonding between two non-metals or non-metal and metal and involves sharing electrons. An example is O2 - oxygen.\n\nFinally hydrogen bonding is self-explanatory consisting of hydrogen and another element.\n\nBecause of the above, it is chemically impossible for two metals to bond as they will not have enough electrons to stabilize each other in the first place. \n\nI even performed this experiment in class. When a metal is heated it "oxidizes". Its atoms fuse with the oxygen atoms in the air and form the gas. For this reason a metal starts turning brown, as you are simply speeding up the process of rusting.\n\nSo your question is kind of impossible from what I have learned in class. \n\nI'm sorry if I am completely of the point in this answer (if there is something I am not getting and that two metals can actually bond) and I apologize for my ignorance and stupidity in advance. haha
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see if this page helps:\n\nhttp://www.chemsoc.org/visElements/pages/data/intro_groupi_data.html
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Because they effectively bind to but also dissociate from the metal center.  They protect the metal from reaction with undesired things until something which is a better ligand comes along and knocks it off.\nIt also serves to generate low-valent or even zero-valent forms of the metal (such as palladium tetrakis).
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Intermetallic bonding is definitely confusing....\n\nFirst to review titanium chemically....All 'transition' metals are called as such because the the electrons in their outer-most shells keep shifting from one orbital to another and never stay put. This is why transition metals form compounds with variable oxidation states/valencies, as they don't always donate a fixed number of electrons, unlike group 1 and 2 elements which always donate 1 and 2 electrons respectively. So if a transition metal like titanium is to combine with another metal, then the resulting compound would not have a fixed formula, as you justified by mentioning TiAl and TiAl3.\n\nWhen 2 metals are to combine, both of them would donate electrons and since neither of the 2 metals would accept them, the donated electrons form an electron cloud around the atoms. Imagine it as 2 positive charges (nuclei of both metallic atoms) surrounded by a negatively charged cloud. The electron cloud is held in place by the attractive force of the nuclei for the electrons. \nIf 2 non-transition metals combine, the electron cloud formed would contain the fixed number of electrons that each metal would donate. If 2 transition metals combine or a transition metal combines with a non-transition metal, the electron cloud would vary in the number of electrons in it. \n\nConsider the electronic configuration of Ti and Al: \n1s2, 2s2, 2p6, 3s2, 3p6, 3d2, 4s2. [titanium]\n1s2, 2s2, 2p6, 3s2, 3p1 [aluminum]\nTitanium would donate 4 electrons from 4s2 and 3d2 to achieve noble gas configuration of argon. As stated at the beginning, the electrons in 3p and 3d orbitals would keep shifting to and fro, so it may happen that at one point, 2 electrons from 3s orbital shift to 3d, and the like. So each time titanium has to bond with aluminum, it may donate 4 or more than 4 electrons.\nIf at one point, 1 titanium atom donates 4 electrons from 4s2 and 3d2 and 4 aluminum atoms donate 1 each from 3p1, then both will have a complete orbital (for Al) and a complete shell (for Ti) left. So with 1 Ti atom and 4 Al atoms, the compound formed will be TiCl4. Other compound like TiAl will be formed according to the number of atoms (of both metals), and the number of electrons they donate.
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I will get you started on the exact answers to this, my friend.  But I am not going to do all your work for you.\n\nGo to http://myweb2.search.yahoo.com/search?fr=ks&p=paramagnetic+elements \n\nOnce you pull that up and start reading, it will lead to all the rest fairly directly.\n\nBe sure to groove on each definition to get the exactitude of it -- magnetic, paramagnetic, diamagnetic, etc.\n\nTry to learn something about the electron configurations involved.\n\nAll good things take some work, you will learn better if you do it yourself while taking notes.