1. Chapter 8 – Basic Chemical Bonding

2. Why are bonds important?
Bonds are what hold atoms together
Without them, water (H2O) would not be water– it would simply be hydrogen and oxygen.
Bonds are the “sticks” that connect the little balls of molecules together, and make everything we see.

3. How do bonds work?
+ Bonding is all about electrons. + Atoms either share or donate electrons to each other in a bond in order to have eight electrons

4. what type of electrons?
+ Bonding involves valence electrons + Why only the valence electrons? + Let’s practice drawing Lewis structures.

5. What determines the type of bond?
Bond type is determined by differences in electronegativity. Electronegativity – The ability of an atom to attract electrons in a bond. In other words how greedy an atom is for electrons.

6. What determines the type of bond?
BIG differences transferring electrons
small differences sharing electrons

7. Ionic Bonding – a bond between a metal and a nonmetal where one atom accepts/gains an electron and the other donates/loses an electron
The more electronegative, nonmetal element ACCEPTS electron(s)
The less electronegative, metal element DONATES electron(s)
This is due to a BIG difference in electronegativity.

8. Example of Ionic Bonding
Chlorine steals sodium’s one valence electron to fill its valence shell.

9. Ionic Bonding
Magnesium + Iodine
Sodium and oxygen
Is Mg happy? I Mg I

10. Isoelectronic
When a nonmetal and a representative-group metal react to form a binary ionic compound, the ions form so that the valence electron configuration of the nonmetal achieves the electron configuration of the next noble gas atom. These ions/atoms containing the same number of electrons.
Ex: F-,Na+, Mg2+, and Al3+ has the same # of electrons just like the noble gas Ne.

11. Ionic Bonds Ionic bonds occur between a cation and an anion.
* The cation is positively (+) charged. (the one that loses the bone)
* The anion is negatively charged. (the one that gets the bone)

12. Properties of Ionic Bonds
Electrons are transferred between a metal and a non-metal.
Ionic solids are crystal solids at room temperature.
High melting points
Good conductors when dissolved.
Have a positive and negative charge.

13. Covalent Bonds Covalent bonds are all about sharing electrons.
Co- meaning “together” or “shared”
-Valent – meaning “outside electrons”
Covalent bonds are all about sharing electrons.
Bonds between atoms with small differences in electronegativity tend to be covalent.
This chart shows differences in electronegativity:
Ionic Polar Covalent Non-Polar

14. Polar Covalent Bonding
Polar bond is a covalent bond with greater electron density around one of the atom in a compound. On the periodic table, electronegativity generally increases across a period and decreases down a group. The range of electronegativity values is from fluorine the most electronegative (4.0) to cesium the least electronegative with a value of 0.7.

15. Example of a Covalent Bond:
In hydrogen gas, which is a diatomic (di = two) molecule, each of hydrogen’s electrons are shared to make H2.
Identical, bonded atoms like H2 have an electronegativity difference of zero.

16. Properties of Covalent Bonds
Covalent bonds happen between atoms on the same side of the periodic table.
They tend to have:
Lower melting and boiling points
Can exist at any state/phase at room temp.
Poor conductivity (electrons can’t move easily)

17. Lewis Structures for Covalent

18. Lewis Structure
Shows how valence electrons are arranged among atoms in a molecule& attains stability by having the noble gas electron configuration.
Hydrogen forms stable molecules where it shares two electrons(Duet Rule).
Most elements form stable molecules by attaining eight electrons (Octet Rule).

19. Covalent Bonds
Single Covalent Bond: A covalent bond in which two atoms share one pair of electrons. Ex: H2, F2 Double Covalent Bond: A covalent bond in which two atoms share two pairs of electrons. Ex: O2 Triple Covalent Bond: A covalent bond in which two atoms share three pairs of electrons. Ex: N2

20. Steps for Writing Lewis Structures
Sum the valence electrons from all the atoms.
If it is an anion, add one electron for each negative charge.
If it is a cation, subtract one electron for each positive charge.
Draw skeletal structure of compound showing what atoms are bonded to each other. Keep the least electronegative element in the center.
Atoms usually have noble gas configurations. Arrange the remaining electrons to satisfy the octet rule or duet rule for hydrogen.
If the structure contains too many electrons, form double or triple bonds on the central atom.

21. Lewis Dot Structure Ex: Draw the Lewis Dot Structures for CCl4 CO2
BCl3
H2O
NH4+

22. Resonance Structures: More than one valid Lewis structure can be written for certain molecules. Ex: NO3–

23. Exception to the octet rule: ( less than the octet) Boron tends to form compounds in which the boron atom has fewer than eight electrons (incomplete octet). BH3 = 6 ve–

24. Exception to the octet rule: ( more than the octet) SF4 = 34 ve–

25. Formal Charge
Formal Charges are used to evaluate the best plausible Lewis structure for a molecule. Atoms in molecules try to achieve formal charges as close to zero as possible. The best structure is the one in which negative formal charges are placed on the more electronegative atoms.

26. Formal Charge
Formal charge = (# valence electrons in free atom) – total # non bonding electrons- 1/2(total # bonding electrons). The sum of the formal charges of the atoms in a molecule or ion must equal the charge on the molecule or ion. Ex: What is the best Lewis structure for CH2O ?

27. ΔH =(bonds broken) – (bonds formed)
Bond Energies
ΔH =(bonds broken) – (bonds formed)

28. Given the following information: Bond Energy (kJ/mol) C–H 413
Ex: Predict ΔH for the following reaction:
CH4 + Cl2  CH3Cl + HCl
Given the following information:
Bond Energy (kJ/mol)
C–H
Cl–Cl
C–Cl
H- Cl
[3(413) ] – [3(413) ] = –42 kJ
ΔH = –42 kJ