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Lewis Dot Structures

Lewis Dot Structures. Lewis. Lewis Dot Structures. For atoms--- Figure out how many valence electrons an atom has. Draw the electrons one at a time around the atoms symbol in a counter clockwise fashion starting on the right side and then going to the top, other side and then the bottom.

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Lewis Dot Structures

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  1. Lewis Dot Structures Lewis

  2. Lewis Dot Structures For atoms--- • Figure out how many valence electrons an atom has. • Draw the electrons one at a time around the atoms symbol in a counter clockwise fashion starting on the right side and then going to the top, other side and then the bottom. • After one electron is on each side, double them up until you run out of electrons.

  3. Ionic CompoundsVsCovalent(Molecules) Compounds

  4. Ionic Compounds • The binary variety is made up of a metal and a non-metal • Have a difference in electronegativity of more than 1.9 (chapter 9 pg. 304) • Formed from ions • Electrons are transferred • Electrostatic attraction!

  5. Covalent Molecules • Made up of two non-metals • Have a difference in electronegativity of less than 1.7 (can use back of periodic table to see) • Ions are not formed • Electrons are shared

  6. Polar vs. nonPolar • Electronegativity difference of .5 to1.9 Polar (electrons in bond are not shared equally) • Electronegativity difference below .5 nonPolar (equal sharing of electrons in bond)

  7. Rules for writing ionic compound Lewis structures • Determine that a metal is present / first • element on left of periodic table(Not counting H). • Write the Lewis structures for the elements in • the ionic compound. • Draw an arrow showing the transfer of the • electrons from the metal (low electronegativity • element) to the non-metal (high electronegativity.)

  8. Ionic Compounds

  9. Molecular Compounds • Have a difference in electronegativity of less than 1.7 • Usually composed of two non-metals

  10. Rules for writing Covalent compound Lewis structures (Molecules) 1. Draw the Lewis structure for each of the elements 2. See how many of the electrons are unpaired and could be shared so that the atoms will each get eight around them or two for H. 3. Circle the shared pairs of electrons. Example NH3

  11. 1. -Sum the valence electrons from all the atoms in the molecule. -Do not worry about which electron comes from which atom. -It is only the total number of electrons that is important.

  12. 2. -Use a pair of electrons in the form of a line to show a bond between each pair of atoms. -The central atom is usually by itself in the formula.

  13. 3. -Arrange the remaining electrons to satisfy the duet rule for Hydrogen and the octet rule for other elements. -Carbon, Nitrogen, Oxygen and Fluorine (CNOF) will never violate the octet rule, other elements can. -If an atom violates the OCTET Rule, it usually is the central atom.

  14. 4. Place any remaining electrons around the central atom in pairs! - again if there are exceptions to the octet rule they will usually be the central atom.

  15. 5. check valencies - Every atom has a certain number of reactive (valence) electrons that it brings to the bonding “table”. When you have finished placing all of the electrons available in bonds or around atoms the atoms should have their own electrons and should be sharing electrons up to 8!

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