Bonding

Lewis Dot Structures 

  • Place valence electrons(dots) around the element symbol for the atom 

  • Place on each side 

Lewis Dot Structures(Compounds)

  • Count up total # valence electrons in the compound

  • Draw atoms → Least electronegative metal in the middle 

  • Connect w/ single line 

  • Check to see what atoms need more e- to be full & draw dot on then in pairs - outside, then middle  

Ionic Bond 

  • Metal and nonmetal(exception: Beryllium)

  • Electrons are transferred from metal(+ always) to nonmetal(- always) forming ions 

  • The strong attractive force between ions of opposite charges(charges must cancel out ot make a neutral compound) 

  • Ions attract object w/ opposite charges & organize into crystal(repeating arrangement of ions)

Ionic Compound

  • Tightly bound

  • Hard, brittle, solid at room temperature, high melting point, dissolve in water 

  • Electrical conductivity needs movement of charged particles. 

Covalent Bonds 

  • Occurs between 2 nonmetals or a nonmetal and metalloid electrons are shared between atoms 

  • An uncharged group of 2 or more atoms held together by covalent bonds is a molecule. 

  • Liquid or gas at room temperature

  • Do not conduct electricity 

  • Do not dissolve in water

Diatomic molecules 

H,O,Br,F,N,Cl

Determining bond type w/ electronegativity

Less than 0.3 → non polar covalent 

0.3 - 1.7 → polar covalent 

More than 1.7 →Ionic  


Resonance Structures 

  • More than one possible correct Lewis Structures for a compound, those structures are called resonance structures.

Formal Charges on Lewis Structures 

  • When drawing Lewis structures for charged covalent molecules(polyatomic ions) you may end up with mult. structures

  1. Polyatomic Ions - a group of atoms that are covalently bonded that have a charge.

  • Formula for formal charges 

  1. Electrons - # of lines - # of dots 

  • Rules for formal charges 

  1. Lowest possible formal charges 

  2. Least amount of charges possible 

  3. Charge located on the more electronegative atom 

  • Molecular Geometry

  1. VSEPR - Valence Shell Electron Pair Regulation

  2. Predicts shade of molecules & ions by assuming that the valence shell pair are arranged as for as possible 

Basic Shapes 

How to tell what compounds go with each shape

  • Linear AB or AB^2(pretend that it’s on the bottom)

  • Trigonal Planar AB^3 

  • Tetrahedral AB^4 

  • Trigonal pyramidal AB^4 (3 bond pairs and 1 lone pair)

  • Trigonal bipyramidal AB^5

  • Octahedral AB^6

How to determine the shape of a molecule 

Number of atoms attached to central atom 

0 lone pairs 

1 lone pair 

2 lone pairs 

2 atoms

Linear with 180 degrees

Bent 120 degree

Bent 109.5 degrees 

3 atoms 

Trigonal Planar 120 degrees

Trigonal Pyramidal 109.5 degrees

4 atoms

Tetrahedral 109.5

5 atoms attached  to center would be trigonal bipyramidal and 6 atoms attached to center would be octahedral 

 Molecular Polarity 

  • Polarity is the result of the uneven distribution of electrons in a molecule, as a result of different electronegativities. 

  • Dipole moments: Show the separation of charge leading to a molecule with a negatively charged end and a positively charged end.

How to determine Polarity 

Some ways to determine polarity are:

  • If there are multiple lone pairs around a common center   

    • ex: H2O

  • If the molecule is asymmetrical ex: CH3Cl or CHCl3

  • If the dipoles do NOT all point to the same, common center

Types of intermolecular forces 

Intermolecular forces - the force of attraction between 2 molecules 

  • Hydrogen bonds

  • Dipole-dipole 

  • London dispersion forces 

Hydrogen Bonds 

  • Attraction between hydrogen and a very electronegative atom (F, O, and N only) 

  • H bonding is FON :D

  • Results in a dipole because of the uneven pull for electrons

  • Ex: H2O molecules, HF, NH3, etc.

  • Strongest attraction

Dipole - Dipole Bonds 

  • Force of attraction between two or more polar molecules ex: PCl3, HF, NH3, etc. 

  • Due to the uneven pull for electrons because of the difference in electronegativity between the two atoms

  • Strong attraction

London Dispersion Bonds 

  • Force of attraction between two or more nonpolar molecules

  • Ex. Present between diatomic molecules such as O2, N2, and H2 in the air

  • Weak attraction

Metallic Bonds 

  • Formed between metal atoms

  • results when metal atoms, from the same element, release their valence electrons to a pool of electrons (“sea of electrons”) shared by all of the metal atoms

  • Ex: copper atoms

  • Most are solid at room temperature

  • Luster

  • malleability

  • ductility

  • good conductor of heat and electricity


Alloys 

  • Mixture made up of two or more metals or metal and carbon and held together by metallic bonding

  • Alloys are often more stable, stronger, and more resistant to corrosion than the individual metal atoms.
    - Examples:

  • stainless steel – 72.8% iron, 17% chromium, 7.1% nickel,

    • ~1% of each aluminum and manganese

sterling silver – 92.5% silver and 7.5 % copper