Unit 3 Chemistry(Polarity, Ionic/Covalent, Lewis structures)

  $$B_4H_{10}$$  

$$B_2H_9$$

$$B_5H_{10}$$

$$B_5H_9$$

-3

-2

+2

+3

N-Cl, P-Cl, As-Cl 

P-Cl, N-Cl, As-Cl

As-Cl, N-Cl, P-Cl

As-Cl, P-Cl, N-Cl

The electrons in a bond between two oxygen atoms (O2) are shared equally, and the resulting bond is polar.

The electrons in a bond between two oxygen atoms (O2) are shared equally, and the resulting bond is nonpolar.

The electrons in a bond between two oxygen atoms (O2) are shared unequally, and the resulting bond is polar

The electrons in a bond between two oxygen atoms (O2) are shared unequally, and the resulting bond is nonpolar.

diamond (C)

copper (Cu)

calcium oxide (CaO)

hydrogen sulfide (H2S)

It is metal.

There are extra electrons in copper that can move freely.

It can absorb photons and then re-emit them

Its electrons are present in a delocalized “sea” around metal nuclei.

Ionic bonds are much stronger than the molecular attractions between covalently bonded compounds. 

Covalent bonds are not as strong as ionic bonds. 

As a solid, ionic compounds have a very organized crystalline structure which takes a lot of energy to break apart.

Most covalent compounds have at least one weak bond in their structure that is easily broken when heat is added.

Network covalent solid because there are covalent bonds between the atoms.

Ionic solid because there are positive and negative charges attracted to each other.

Molecular compounds because metals and nonmetals are sharing electrons.

Metallic solid because electrons are distributed among the positive nuclei.