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AP Chemistry 12
Predicting Properties Based on Bonding
Predic'ng Proper'es Based on Bonding
• So far, we have been discussing bonding within a molecule, this type of interaction is known as intramolecular
• Intramolecular interactions don’t really tell us anything about the properties though… For Example: v Why is H2O a liquid at room temperature, but H2S and H2Se are gases? (Both are bent geometry)
v Why is I2 a poor electrical conductor but graphite (C) a great conductor, even though both are non-‐metals?
Predic'ng Proper'es Based on Bonding Bonding
Ionic
Lattice Bonding
Metallic
Delocalized Electrons
Covalent
Molecular Covalent
IMF (London Dispersion, Dipole-‐Dipole Interactions, Hydrogen Bonding)
Network Covalent
Extended Covalently Bonded Structure
Ionic Bonding
• In ionic compounds, the metal loses electrons and the non-‐metal gains them • The metal becomes positively charged and the non-‐metal becomes negatively charged • The positively charged ion (cation) and the negatively charged ion (anion) attract one another
Reacting Elements:
Electron ConViguration:
During Reaction:
New Electron Arrangement:
Ions Formed:
Na Cl
1s22s22p63s1 1s22s22p63s23p5
loses 1e-‐ gains 1e-‐
1s22s22p6 1s22s22p63s23p6
transfer of an electron
e-‐
Na Cl -‐ +
Electrosta'c Attraction
Ionic compounds form a LATTICE STRUCTURE. Millions of oppositely charged ions are held together in a very stable arrangement.
-‐ Na
+ Cl
Electrosta'c A5rac'on
• The strength (lattice energy) of this attraction is given by the formula:
Q = the charge on the ions r = the distance between them • This formula shows that higher charges and smaller radii ions have the strongest attraction
• Exact formula is called Coulomb’s law
E∝Q1Q2
r
La7ce Energy
• We don’t need to actually calculate the lattice energy in this case, but simply compare the energies qualitatively between different ionic compounds • Example: LiF vs LiI vs LiF • Example: NaF vs MgO
LiF vs LiI vs LiF and NaF vs MgO
Physical Proper.es of Ionic Solids
1. Brittle, Crystalline, Hard 2. High Boiling and Melting Points (Strong
attraction) 3. Conductive when melted or dissolved in
water (electrolytes). Any time ions are formed, the solution will conduct electricity. (l or aq)
4. Non-‐conductive as solids or gases (s or g) 5. Low vapour pressure 6. Generally soluble in water (many
exceptions; recall solubility table)
Side Note on Vapour Pressure • Molecules in the liquid or solid phase can escape into the gas phase due to random motion (vapourization)
• At the same time, molecules in the gas phase will condense back into the liquid or solid phase (condensation)
• When both rates are equal, the number of molecules in the gas phase remain constant, we can measure the pressure of this gas. This is called the vapour pressure.
Side Note on Vapour Pressure • Vapour pressure depends on the temperature of the system. Vapour pressure increases with temperature. v More molecules have enough energy to escape and stay in the gas phase
• Strong attractive forces between molecules result in lower vapour pressure while weak attractive forces result in higher vapour pressure (indirect proportionality)
• Substances with high vapour pressure are known as volatile substances
Predic'ng Proper'es Based on Bonding Bonding
Ionic
Lattice Bonding
Metallic
Delocalized Electrons
Covalent
Molecular Covalent
IMF (London Disperson, Dipole-‐Dipole Interactions, Hydrogen Bonding)
Network Covalent
Extended Covalently Bonded Structure
Metallic Bonding • Metal atoms share its valence electrons with its neighbours • The positive metal nucleus are embedded in many valence electrons • The valence electrons are free to roam, acting similar to a “sea of electrons”
Metallic Bonding
Physical Proper'es of Metallic Solids
1. Malleable, Ductile, Strong 2. Generally high melting point 3. Conductive as solids and liquids 4. Low vapour pressure 5. Not soluble in water
Alloying • Because many metal cations are similar in size, small amounts of another metal can be added • When atoms of similar sizes are used to replace metal atoms, the result is a substitutional alloy • An example of this is brass
v Copper and Zinc, both have 0.135 nm radius
Alloying • Because there are small spaces between metal atoms, other atoms can Vit in between them when alloying • When atoms of smaller radii are used to alloy, a interstitial alloy is created • An example of this is steel
v Iron and carbon, 0.125 nm and 0.070 nm radii
Pure metal: Iron, Fe
Substitutional Alloy: Brass, Cu/Zn
Interstitial Alloy: Steel, Fe/C
Mixed Alloy: Stainless Steel, Fe/Cr/C
Alloying
Drawing Alloy Prac'ce
• Substitutional Alloy v 50 % Copper and 50 % Zinc by mass (Brass) v Convert the mass to nearest whole # mole ratio v Draw at least 30 atoms in total
• Interstitial Alloy v 90 % Iron and 10 % Carbon by mass (Steel) v Draw at least 30 atoms in total
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