1. Solids
Chem 112

2. Solids: General Info. Atoms vibrate Kinetic energy- low
Diffusion- extremely slow
Evaporation- usually slow, except substances that sublime
Sublimation: transition of solid directly to vapor.
Example: Ice cubes slowly “disappear” in the freezer.

3. Bonding in Solids
In crystalline solids atoms are arranged in a very regular pattern.
Amorphous solids are characterized by a distinct lack of order in the arrangement of atoms.

4. Often considered to be super cooled liquids
Amorphous materials- Not true solids, they have no distinct crystal lattice (repeating pattern)
Often considered to be super cooled liquids
Examples: peanut butter, wax, clay, silly putty, and glass

5. Although glass appears to be a solid, it isn't
Although glass appears to be a solid, it isn't. The molecules are connected in no special way. Solids have more of an ordered molecular structure.

6. Bonding in Solids There are four general types of solids.
Metallic solids share a network of highly delocalized electrons.
Ionic solids are sets of cations and anions mutually attracted to one another.

7. Bonding in Solids
Covalent-network solids are joined by an extensive network of covalent bonds.
Molecular solids are discrete molecules that are linked to one another only by van der Waals forces.

8. Molecular Solids
Held together by Van der Waals forces (covalent bonds, LDF, dipole-dipole and hydrogen bonds)
Soft, low to moderately high melting points
Poor conductors
Ex. SucroseC12H22O11

9. Molecular Solids
Graphite is an example of a molecular solid, in which atoms are held together with van der Waals forces.
They tend to be softer and have lower melting points.

10. Covalent-Network Solids
Diamonds are an example of a covalent- network solid, in which atoms are covalently bonded to each other.
They tend to be hard and have high melting points, poor electrical conductors

11. Metallic Bonding/Metallic Solids
In elemental samples of nonmetals and metalloids, atoms generally bond to each other covalently.
Metals have plenty of valence electrons and form large groups of atoms that share electrons among them.

12. Metallic Bonding/Metallic Solids
You can think of a group of cations suspended in a sea of electrons.
The electrical and thermal conductivity, ductility, and malleability of metals is explained by this model.

13. Metallic Solids Held together by metallic bonds Soft to hard
Low to very high melting points
Good conductors
Malleable and ductile
Gold

14. Crystal Lattices
we can find the pattern in a crystalline solid by thinking of the substance as a lattice of repeating shapes formed by the atoms in the crystal.

15. Crystal Lattices
The individual shapes of the lattice, then, form "tiles," or unit cells, that must fill the entire space of the substance.

16. Crystal Lattices There are seven basic three-dimensional lattices:
Cubic
Tetragonal
Orthorhombic
Rhombohedral
Hexagonal
Monoclinic
Triclinic

17. Crystal Lattices
Within each major lattice type, additional types are generated by placing lattice points in the center of the unit cell or on the faces of the unit cell.

19. Cubic Structures
we can determine how many atoms are within each unit cell which lattice points the atoms occupy.

20. Close Packing
The atoms in a crystal pack as close together as they can based on the respective sizes of the atoms.

21. Hexagonal close packed

23. lattice structure

24. Alloys
In substitutional alloys, a second element takes the place of a metal atom.
In interstitial alloys, a second element fills a space in the lattice of metal atoms.

25. Ionic Solids
In ionic solids, the lattice comprises alternately charged ions.
Ionic solids have very high melting and boiling points and are crystals.

26. Ionic Solids
Form crystal lattices- the particles are arranged in a repeating pattern
diamond
galena
NaCl
Ice cube

27. Chemical Bonds Three basic types of bonds Ionic Covalent Metallic
Electrostatic attraction between ions.
Covalent
Sharing of electrons.
Metallic
Metal atoms bonded to several other atoms.

28. Lattice Energy Eel =  Q1Q2 d lattice energy:
The energy required to completely separate a mole of a solid ionic compound into its gaseous ions.
The energy associated with electrostatic interactions is governed by Coulomb’s law:
Eel = 
Q1Q2 d

29. Lattice Energy
Lattice energy, then, increases with the charge on the ions.
It also increases with decreasing size of ions.

30. Example:
Use the periodic table to arrange the ionic compounds NaF, CsI, and CaO in order of increasing lattice energy
You know that the larger the ionic charges the greater the energy, and the farther apart the ions are, the lower the energy
SO..

31. Example continued… Q1Q2 Eel =  d
1. Check magnitude (size) of charge first
2. Difference in distance between ions
Ionic size increases going down a group so we know that Cs+ is larger than Na+
We also know that I -1 is larger than F -1
That makes the distance between Na+ and F- smaller than the distance between Cs+ and I-
Lattice energy is greater when the distance between ions is smaller
NaF = Na +1 and F-1
CsI = Cs + and I -1
CaO = Ca+2 and O -2
Based on charge we expect CaO to have the largest lattice energy
Charges are the same for NaF and CsI so go to step 2 to check distance

32. Example continued … Final answer:
In order of increasing lattice energy…
CsI < NaF < CaO
***The higher the lattice energy, the higher the melting point***

33. Energy changes and lattice energy of ionic compounds
Lattice energy – is a measure of the strength of ion-ion interaction
- determined by the change in energy when free ions in the gas
phase combine to form 1 mole of solid ionic compound