1. CHEMICAL BONDING

2. ***Occurs when atoms of elements combine together to form compounds.*****

3. Formation of compounds
Involves valence electrons.
PE is lower in bonded atoms.
Attractive force that develops is called "chemical bond“
Occurs during chem. reactions

4. Two (or three) methods Ionic bonding - attraction of ions
Covalent bonding - shared pairs of electrons
Metallic bonding - alloys (metals) not compounds

5. An ionic bond -(electrovalent)
a. Definition -  An attraction that forms between oppositely charged ions
b.  Pos. ions + Neg. ions → neutral compound
∆EN > 1.7

6. Ionic Bonds
Cations: Metals lose electron(s) and become positive ions(+).
Examples: Na+1, K+1, Mg+2, Ca+2, Al+3
Anions: Nonmetals gain electron(s) & becomes negative ions(-).
Name ends in “–ide”
Examples: Cl1-, Br1-, O2-
Choride, bromide, oxide . . .

7. Ionic bonds - (crystals)

8. Characteristics of Ionic Solids
Made from + and - ions
Metal-nonmetal (or polyatomic ions)
Compound is neutral.
Tend to be solid,
Brittle and crystalline
High MP and BP
Have strong attractions in all directions
Non-conductors as solids, but will conduct when molten or dissolved
Some dissolve readily

9. Ionic solids are brittle
due to Crystal lattice + -

10. Ionic solids are brittle
Strong Repulsion breaks crystal apart. + - + - + - + -

11. Lattice Energy (ionic)
More correct than bond energy for crystals.

12. Metallic Bonding Sea of Electrons
In metals, the valence electrons are not bonded to any specific atom. (delocalized)
Able to move freely over the positive centers.
Causes unique properties.

13. Characteristics of Metals
A. Malleable: dent when hammered.
B. Ductile: draw into a wire
C. Conductivity: electricity and heat
D. Alloy: a blend of metals

14. Sea of Electrons + Electrons are free to move through the solid.
Metals conduct electricity and heat well. +

15. Malleable +

16. Malleable
Electrons allow atoms to slide by. + + + + + + + + + + + +

17. Alloys Alloys: mixtures of two or more metals.
Important because their properties are often better than the individual elements.
Examples:
Bronze is made from copper & tin. It is harder than copper & more easily cast.
Sterling silver: Ag (92.5%) & Cu (7.5%)
Stainless steel: Fe (80.6%), Cr (18%), C (0.4%), & Ni (1%)
-- Brass, Pewter, and others.

18. Covalent Bond
a. an attractive force that develops between atoms that are sharing pairs of electrons
∆EN < 1.7
b. Hydrogen – H2
H• + H• → H:H (dot diagram)
Structural formulas use a dash H - H

19. Characteristics of Covalent Compounds
Nonmetal-nonmetal combinations
Can be gases, liquids, or solids
Low to med. MP and BP
Insulators/Nonconductors (except for acids)
Molecular (a few are crystalline)
Generally not soluble (some polar exceptions)

20. Bond energy Energy required to break a bond.
Bonds form to lower PE, so breaking bonds will increase PE.

21. Breaking Bonds is always endothermic
Energy is required.
ALWAYS.

23. Covalent Bonds Why not ionic?
The difference in electronegativity is less than 1.7
Electrons are not pulled away from either atom.
They are shared.

24. Fluorine – F2

25. Ammonia -NH3

26. Special Types of Covalent Bonds
Multiple bonds – Occasionally atoms share more than one pair of electrons
Double bond – two shared pairs
Ex. O O=O or

27. - Triple bonds - two atoms share three pairs of electrons. Ex. N2
N N or

28. How to write Lewis Structures
1. Set out the atoms – think symmetry.
2. Count all the valence e-
3. Insert single bonds first, then fill rest.
All the e- are paired.
each nonmetal atom requires an octet.
H only requires 2 e-.
Multiple bonds may be needed.
Readily formed by C, N, O, S, and P.

29. Polarity
An unequal sharing of electrons due to difference in electronegativity.
Polar bond – Any bond with ∆EN
Polar molecule – Has a positive end and a negative end.
Occurs in water and ammonia (**know these two)
Causes intermolecular attraction to increase

30. Molecular Substances
Covalently bonded substances – show more variety in phases and properties. Tend to be insulators(nonconductors).
Nonmetal-nonmetal combinations
Nonpolar – tends to be gases at room temp. – have only dispersion (Vanderwaals) forces. Have low MP and BP and high VP(vapor pressure)
Polar –(dipoles) tend to be liquids or solids at room temp. Have ↑MP and BP and↓VP
H- bonding – very strong type of polar

31. VALENCE SHELL ELECTRON PAIR REPULSION
VSEPR
VALENCE SHELL ELECTRON PAIR REPULSION
A theory which describes the shapes of molecules based on the idea that pairs of electrons will repel each other as much as possible.

32. Polar molecules are asymmetrical and have polar bonds.
Nonpolar molecules are symmetrical and may or may not have polar bonds.

33. Intermolecular Attractions (Vanderwaals forces)
Attractive forces between molecules
Happens to covalent molecules.
Strongest - hydrogen bonds
Medium dipoles/ polar
Weakest - dispersion or London forces
( All molecules have London forces)

34. Hydrogen bonding
In some highly polar compounds, a H atom is attracted to, and forms a weak bond with, an adjacent molecule.
Only occurs in compounds where there is:
H-F (strongest)
H-O
H-N (weakest)
(know these three)

35. Hydrogen Bonding H O d+ d- H O d+ d-

36. Hydrogen bonding H O H O H O H O H O H O H O

37. H-bonding Gives water its very unusual properties. High MP and BP
Holds the DNA molecule together.
Provides stability and shapes for proteins, enzymes, etc.
Strongest type of intermolecular attraction.

38. London Dispersion Forces(Van der waals)
Weakest type of intermolecular attraction.
Develops between nonpolar molecules due to temporary shifts in the electron positions.
Strength of attraction is directly proportional to the number of electrons
(wax is a nonpolar molecule – but large so it is solid)

39. Network Solids Large arrays of covalently bonded crystals.
Do not conduct, hardest solids
Very high MP and BP
Examples: diamond, graphite, SiO2 (very few – easiest to just memorize)
More on this in 2nd semester

40. Bond length and energy
Bond length depends on the size of the atoms/ions and the number of bonds between them
C-C is longer than C=C is longer than CC
Shorter bonds are stronger.
Measured in kJ/mol

41. Breaking Bonds is always endothermic
Energy is required.
ALWAYS.

42. Endothermic reactions

43. Forming bonds is exothermic
Energy is released.
ALWAYS

44. Exothermic reactions