1.2. In metallic bonds, the valence electrons from the s and p orbitals of the interacting metal atoms delocalize.
1.3. Metals are lustrous, malleable, ductile, good conductors of heat and electricity
1.4. Metals tend to have high melting points and boiling points suggesting strong bonds between the atoms.
1.5. Metals are insoluble in water or organic solvents unless they undergo a reaction with them. Typically this is an oxidation reaction that steals the metal atoms of their outer electrons, destroying the metallic bonding.
1.6. Metallic bonds hold the metallic solid together. Atoms are arranged like closely packed spheres. Because outer electrons of metal atoms are delocalized and highly mobile, metals have electrical and thermal conductivity.
1.7. Lewis Dot Structures:
2. Created By: Mitchell Meister
3. Ionic
3.1. Ionic compounds generally form between elements that are metals and elements that are nonmetals.
3.2. Ionic bonding is the complete transfer of valence electron between atoms. It is a type of chemical bond that generates two oppositely charged ions
3.3. Solid; High Polarity
3.4. Ionic solids typically melt at high temperatures and boil at even higher temperatures
3.5. Polar compounds tend to dissolve in water, and we can extend that generality to the most polar compounds of all ionic compounds
3.6. Ionic compounds conduct electricity when liquid or dissolved in water, because their ions are free to move from place to place
3.7. Lewis Dot Structures:
4. Covalent
4.1. Covalent bonds usually occur between nonmetals
4.2. Covalent bonding occurs when pairs of electrons are shared by atoms.
4.3. Covalent compounds may exist as a solid, a liquid, or a gas
4.4. They have lower melting points and electrical conductivity compared to ionic compounds.
4.5. Covalent compounds aren't usually very soluble in water because they have a tendency to dissociate or ionize in water
4.6. Solid ionic compounds do not conduct electricity because there are no free ions or electrons
