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Bonding

Bonding. How Atoms Combine and Bond. Underlying principle of all bonding:. Atoms always seek noble gas structure to become stable. They give, take, or share electrons to become stable like a noble gas. Valence Electrons:.

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Bonding

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  1. Bonding • How Atoms Combine and Bond

  2. Underlying principle of all bonding: • Atoms always seek noble gas structure to become stable. • They give, take, or share electrons to become stable like a noble gas.

  3. Valence Electrons: The electrons available to be lost, gained, or shared in the formation of chemical compounds

  4. Examples: • A sodium atom, Na, is unstable with 11 e- • If it loses 1 e-, it becomes stable like Neon, with 10 e- • A Cl atom is unstable with 17 e- • If it gains one e-, it becomes stable like Argon with 18 e-

  5. Octet Rule: Atoms tend to bond with other atoms until they have eight electrons in their outermost electron shell

  6. Exceptions to Octet Rule: Hydrogen only 2 valence e- Li and Be  bond to get 2 And 4 e- B bond to get 6 e- Beyond second row elements can expand into d orbitals and acquire more than 8 e-

  7. Types of Bonds • Covalent (aka: Molecular) • A. Electrons are shared • B. Occurs between non-metals • C. Represented with Lewis Dot Diagrams/Structures D. Forms molecules with molecular formulas – H2O, NH3

  8. Types of Bonds • Covalent (aka: Molecular Cont.) • F. Forming Polar Bonds • 1) The electrons are not shared evenly • 2) Electrons spend more time on the more electronegative atoms 3) If the polar bonds are not balanced the whole molecule is polar!

  9. How to determine how polar a bond is: • Look at how far apart the two elements are across the periodic table; the farther apart they are, the more polar the bond. • List in order from most polar to least polar: H2O, NO2, NaF, Cl2 • NaF is the most polar because it’s a metal and non-metal • H2O is next. Both non-metals, but far apart. • Then comes NO2, close together, so not too polar • Cl2 is totally non-polar • More Precise: Subtract their “electronegativity” values – see page 161 in the text.

  10. Non-polar to ionic:

  11. Remember the Diatomic Elements! • Molecules containing only 2 atoms • Long Live Old Mrs. HOFBrINCl !

  12. Types of Bonds • Covalent (aka: Molecular Cont.) • E. Naming • 1) Drop the last or two last syllables • of the 2nd element and add “ide” • Bromine  Bromide • Nitrogen  Nitride • 2) Add Greek prefixes for the number of atoms (pg. 228)

  13. Practice • CO • carbon monoxide • b. CO2 • carbon dioxide • S4N2 • tetrasulfurdinitride • d. N2O5 • dinitrogenpentoxide

  14. II. Ionic Bonds A. Non-metals steal electrons from metals B. Non-metals and metals acquire opposite charges C. Opposites attract – the compound is held together by electrostatic forces D. Form repeating crystal lattices

  15. II. Ionic Bonds E. Naming 1) DON’T USE GREEK 2) USE Naming sheet 3) Cation 1st (+), Anion 2nd (-)

  16. Practice a. BaF2 Barium Floride b. FeBO3 Iron (III) Borate 4) In Reverse - Balance Charge! a. Potassium Chromate K2CrO4 b. Tin (IV) Bromate Sn(BrO3)4

  17. III. Metallic Bonding A. Free-floating valence electrons are shared by all positively charged nuclei - a “Sea of Electrons” B. Properties 1) High electrical and thermal conductivity 2) High luster- due to ability to absorb a wide range of light frequencies.

  18. B. Properties (Cont.) 3) Malleability- the ability to be hammered into sheets. 4) Ductility- the ability to be drawn in to a wire. Malleability & Ductility result from the atoms in one plane of a metal being able to slide past the atoms of another!

  19. Three chemical bonds: • 1. Ionic Bond (metal with non-metal) • 2. Covalent Bond (non-metals) • 3. Metallic Bond (metals)

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