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Chemistry Bonding Covalent Bonding

Chemistry Bonding Covalent Bonding. Do Now: Take out HW to check CALCULATORS NOT REQUIRED PERIODIC TABLES REQUIRED. Covalent bonding is often thought of as “sharing” of electrons between two atoms.

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Chemistry Bonding Covalent Bonding

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  1. ChemistryBondingCovalent Bonding Do Now: Take out HW to check CALCULATORS NOT REQUIRED PERIODIC TABLES REQUIRED

  2. Covalent bonding is often thought of as “sharing” of electrons between two atoms. This occurs when neither atom is strong enough (electronegativity) to forcefully remove electrons from the other, like an atomic tug of war. This “sharing” can result in the appearance of a stable octet for both atoms. Covalent Bonding

  3. The “gaps” in the Lewis dot structure tell us how many bonds that atom is likely to form. 1 bond – halogens and hydrogen 2 bonds – chalcogens 3 bonds – nitrogen group 4 bonds – carbon group I say “likely” because thereare many exceptions tothis… so how do you know? Covalent Bonds by Atom

  4. There is a process for drawing most covalent molecules and ions: Step 1: Count all of the valence electrons and any effects from ionic charges cations: remove electrons, anions: add Step 2: Count all of the electrons all atoms would need to have a stable octet (duet for hydrogen) Step 3: Subtract these two numbers and divide by two. Step 4: The resulting number is the total number of bonds you have to draw in the compound. The least electronegative atom in the formula is usually the one in the center (usually the first listed atom that isn’t hydrogen) THE NUMBER OF VALENCE ELECTRONS SHOULD BE THE SAME BEFORE AND AFTER!! Drawing Covalent Structures

  5. In-Class Examples: molecular bromine (Br2) Bonds are represented with lines or dots. Lone pairs (also unshared pairs) are only represented with dots. There are 7 diatomic elements like bromine above: (elements so reactive that they would rather bond with themselves than stay alone): Br2, I2, N2, Cl2, H2, O2, F2 BrINClHOF Drawing Covalent Structures

  6. In-Class Examples: ammonia (NH3) Bonds are represented with lines or dots. Lone pairs (also unshared pairs) are only represented with dots. Drawing Covalent Structures

  7. In-Class Examples: carbon dioxide (CO2) Molecules can have any combination of single, double, and triple bonds. Due to geometry reasons, there is no such thing as a quadruple bond. Drawing Covalent Structures

  8. In-Class Examples: carbonate ion (CO32-) Ions always are presented in brackets with the charge in the upper right hand corner. This atom will have “formal charges” on its atoms that don’t have the “correct” number of bonds. Drawing Covalent Structures

  9. A formal charge is the difference between the number of electrons an atom normally has ownership over and the amount it is given in a structure. Formal Charge = # of valence electrons – (# of unpaired electrons + # of bonds) Example: Oxygen on carbonate (normally 6 valence electrons) Double bond: 6 VE – (4 unpaired electrons – 2 bonds) = 0 Single bonds: 6 VE – (6 unpaired electrons – 1 bond) = -1 The sum of the formal charges gives the charge on the ion/molecule. Formal Charges

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