Bonding

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Information about Bonding
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Published on February 16, 2008

Author: Melinda

Source: authorstream.com

Bonding :  Bonding All chemical bonds are formed as a result of the simultaneous attraction of two or more electrons. Each atom is trying to achieve the electron configuration of a noble gas (creating filled s and p orbitals which results in a symmetrical arrangement of electrons that cause a stable compound) Valence Electrons:  Valence Electrons Electrons are divided between core and valence electrons. Na 1s2 2s2 2p6 3s1 Core = [Ne] and valence = 3s1 Br [Ar] 3d10 4s2 4p5 Core = [Ar] 3d10 and valence = 4s2 4p5 Valence Electrons:  Valence Electrons 1A 2A 3A 4A 5A 6A 7A 8A Number of valence electrons is equal to Group number. Slide4:  The degree of attraction is based on the difference in their electronegativity values or electron attracting power of the atoms. There are three general classifications of chemical bonds that can be formed. Bonding:  Bonding Three types of bonds 1. Ionic: a bond formed due to the electrostatic attraction between a positive and negative ion caused by an exchange of electrons. 2. Covalent: a bond formed between two non-metals due to a sharing of electrons. 3. Metallic : a bond between metal ions formed due to a “sea of electrons”. Ionic Bonds :  Ionic Bonds These bonds are formed when the electronegativity differences between the two atoms are large (>2.1 or = 2.1) In the tug of war for the electrons, this means that one atom is capable of pulling the electron completely out of the other atom’s energy level and into its own creating positively and negatively charged ions. Ionic Bonding:  Ionic Bonding Ionic bonding: results from the electrostatic attraction between cations and anions. Formation of an ionic bond can be viewed as a transfer of electrons. Slide8:  This causes the formation of a positive ion, cation, and a negative ion, anion. This causes an electrostatic attraction between the positive and negative ion The ionic compound is then composed of alternating positive and negative ions arranged in a crystal lattice structure. Slide9:  The metals would release electrons and become positively charged. The elements which would release the electrons would be from Groups 1 (alkali metals) and Group 2 (alkaline earth metals) along with some of the heavier metals in Groups 13, 14, and 15 and the transition elements. Slide10:  The non-metals would accept electrons and become negatively charged. These elements would be found in the Group 16 ( the oxygen family) and Groups 17 ( the halogens) Properties of Ionic Solids:  Properties of Ionic Solids Crystalline solids at room temperature Have higher melting and boiling points as compared to covalent compounds Most are soluble in water but not soluble in non-polar solvents. Conduct electrical current in molten or solution state Are extremely polar bonds. Naming and Formulas:  Naming and Formulas When naming the ionic solid the metals name always comes first The non-metal comes last but the ending of the non-metal name is dropped and changed to ide. Examples: NaCl = sodium chloride H2S = hydrogen sulfide Li2O = lithium oxide Slide13:  Ionic solids using transition elements: Again start with the metal, but include in parenthesis with Roman numerals the size of the charge, copper(II) or iron (III). Then end with the non-metals ide ending. Examples: FeCl3 is iron (III) chloride SnI4 is tin (IV) iodide CuO is copper (II) oxide Slide14:  Polyatomic ions: The only positive polyatomic ion is ammonium. This would begin the name and the end would follow the rules previously stated for ionic compounds. The negative polyatomic ions each have a specific name and so the metal would follow the previous rules followed by the name of the polyatomic ion. Polyatomic Ions :  Polyatomic Ions The following list are some of the most common polyatomic ions that would be good to memorize. NH4+ ammonium (the only positive ion) Anions: (negative 1 ions) NO3- nitrate NO2- nitrite OH- hydroxide CN- cyanide C2H3O2- acetate MnO4- permanganate Slide16:  Anions (negative 2 and negative 3 ions) SO42- sulfate SO32- sulfite CrO42- chromate Cr2O72- dichromate CO32- carbonate PO43- phosphate Formulas:  Formulas Each formula must end up being electrically neutral. After identifying the charge on each of the positive and negative ions, use subscripts to denote how many of each ion must be used to make the compound electrically neutral. Examples Lithium is 1+ and oxygen is 2- thus you would need two lithiums with one oxygen Li2O Calcium is 2+ and chlorine is 1- thus you need two chlorines for the one calcium, CaCl2 Iron can be 3+ and fluorine is 1- thus FeF3 Iron can also be 2+ and fluorine 1- , thus FeF2 Slide18:  For compounds with polyatomic ions. Ammonium chloride it would be NH4Cl since NH4 is 1+ and Cl is 1- Calcium carbonate is CaCO3 since Ca is 2+ and CO3 is 2- Magnesium phosphate is Mg3(PO4)2 Slide19:  Examples: Sn(OH)2 tin (II) hydroxide Pb(C2H3O2)4 lead (IV) acetate NH4NO3 ammonium nitrate CoSO3 cobalt (II) sulfite Fe2(CO3)3 iron (III) carbonate Ni3(PO4)2 nickel (II) phosphate Covalent bonds:  Covalent bonds Polar covalent bonds are formed when the electronegativity differences between the two atoms are greater or equal to 0.4 and less than 2.1 In the tug of war for the electrons, this means that neither atom is able to completely pull the electron from the other atoms energy levels resulting in an unequal sharing of electrons. Slide21:  Non-polar covalent bonds are formed when the electronegativity differences between the two atoms are less than 0.4 In the tug of war for the electrons, this means that both atoms have virtually the same strength of attraction for the electrons resulting in the electrons being located halfway between both nuclei. Slide22:  In covalent bonded molecules, the charges on the ions are determined by which atom has the greatest electronegativity value. That atom would take it’s normal oxidation state or charge, forcing the other atom to take a more unnatural charge. Slide23:  Examples: CO2 for carbon dioxide, the oxygen is more electronegative resulting in the oxygen taking its normal 2- charge. This forces the carbon to take a 4+ charge. CO for carbon monoxide the oxygen will still take the 2- charge but that forces the carbon to become a 2+ charge. P2S5 Sulfur is more electronegative thus it will take a 2- charge resulting in a total of 10 negative charges. This makes each phosphorus a 5+ charge to create a neutral compound. Naming Covalent Compounds:  Naming Covalent Compounds When forming covalent compounds with two non-metal the use of prefixes is needed. 1 mono 6 hexa 2 di 7 hepta 3 tri 8 octa 4 tetra 9 nona 5 penta 10 deca Slide25:  Examples: CO carbon monoxide CO2 carbon dioxide P2S5 diphosphorus pentasulfide N2O4 dinitrogen tetraoxide

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