Unit 2 AP Chem Vocab

Covalent Bonds

• 2 NONMETALS

• share electrons.

Non polar Covalent Bonds

• Sharing electrons equally

• Has a electron negativity difference of (<0.5) 

Polar Covalent Bond

• Share electrons unequally

• One atom is hogging the electrons

• Electron negativity difference is (0.5<) 

• The atom with the greater electron negativity is hogging the electrons

High potential energy 

• When the atoms are far apart

• The period before they bond together

Low potential energy

• When the atoms are close together

• Bonding

Ionic Bonds

• METAL + NONMETAL bond 

• The metal gives the electron away to the non metal

• HIGHER melting point compared to covalent bonds

Metallic Bonding 

• Delocalized electrons (electrons are able to move freely within the structure

• Good at conducting electricity

Covalent Bonding Graph —> Cl2

• The bond length is where the distance between the 2 atoms is at its lowest

• Bond length (I D) → 200

• Bond Enthalpy (E) → 250 Kj/mole (ALWAYS POSITIVE)

Ionic Bonds/compounds —> Na + Cl

• Na = 1s2 2s2 2p6 3s1 

• Cl = 1s2 2s2 2p6 3s2 3p5

• Na gives one electron to Cl. Cl becomes a gains a octet → Na+ and Cl- →NaCl

• This works because when the metal donates to the non metal, the opposite charges stick together because of their strong electrostatic attractions
  

Melting point in Ionic Bonds

• Melting point increases as the magnitude of the charge increases

• Because of Coulomb's law → LARGER charge = LARGER bond energy = Greater melting point

• EX → Mg+2 and S-2 is stronger than Na+ Cl-

  
  

What if they are the same charge? (Melting points of ionic compounds) 

• melting point DECREASES as ionic size INCREASES. 

• When the ionic nuclei are far apart (due to size), the force of attraction gets weaker. 

  
  

When predicting melting points you should…….

• Look at the greater charge

• Compare sizes (smaller Bonds = higher melting point)

Lattice energy

• energy released when ions are combined into an ionic compound

• STRONGER ionic forces (Based on size and charge) =  Higher amounts of lattice energy released.

Lattice energy examples 

LiF (small) → -1030 Kj/mol 

NaCl (Medium)→ -746 kj/mol → Lower because there is not as much attraction compared to LiF

KBr (Large)→ -688 kJ/mol

Ionic compounds

• High melting point

• Brittle (not malleable) (think salt NaCl)

• Conducts electricity when dissolved in water (electrolytes) 

• Soluble in water (polar substances)

• Has a repeating crystal structure (crystal lattice) 

• Compact — Not floating around

Crystal lattice

• Repeating crystal pattern

• Held together by electrostatic forces

Substitutional Alloys

• Brass, Bronze, etc

• Substitutes into the spot of other elements.

Interstitial Alloys

• EX → Steel 

• Atoms that stick in between the small spaces of the main metal (hardening agent) 

• Allows them to stay in place making the compound stronger

Octet Rule 

• Most atoms are stable with 8 valence electrons

Exceptions to the Octet Rule

• H is stable with 2 electrons 

• B is more stable with 6 electrons

• Non metals with 3 or more energy levels can sometimes have more than 8 valence electrons (up to 10 or 12) → Expanded Octet. 

• When a central atom is bonded to more than 4 atoms, the central atom will have an expanded octet 

Resonance structures

2 lewis electron dot structures which both represent correct structures for a molecule 

Formal charge

• A Way of assigning a charge to each individual atom in a molecule. To have it neutral, each should have a formal charge of 0. 

• Each dot is one 

• A bond gives 1 for each atom it is bonded to 
  

Sigma Bond

Result of overlapping s orbitals

Pi Bonds

Result of overlapping p orbitals  

Single bonds

Longest / Weakest → Sigma bond

Double Bonds

Middle → 1 sigma 1 pi

Triple Bonds

Shortest/ Strongest → 1 sigma 2 pi

Hybridization 

• Largest sum → sp3d2

• Add the number of sigma bonds to the number of unshared electron pairs on the atom 

• Between 2-6

Tetrahedral

• 4 sigma bonds

• 109.5 degrees

Trigonal Pyramidal

• 3 sigma bonds + 1 lone pair

• 107 degrees

•  The bond angle is smaller because the electron pairs exerts more repulsion than shared pairs 

Bent

• 2 sigma bonds + 2 lone pairs

• 105 degrees 

Trigonal planer

• 3 sigma bonds

• 120 degrees

Angular

• 2 sigma bonds + 1 lone pair

• 117 degrees

Linear

• 2 sigma bonds

• 180 degrees

Octahedral

• 6 sigma bonds

•  90 degrees

Square pyramidal

• 5 sigma bonds + 1 lone pair

• 90 degrees

Square Planar

• 4 sigma bonds + 2 lone pairs

• 90 degrees

T shaped

• 3 sigma bonds + 3 lone pairs

• 90 degrees 

Trigonal Bipyramidal

• 5 sigma bonds

• 90 + 120 degrees

See-saw

• 4 sigma bonds + 1 lone pair

• 87 + 117 degrees

Trigonal Planer 

• 3 sigma bonds + 2 lone pairs

• 120 degrees