1. Introduction to Bonding & Chemical Formulas
Chapter 6
Introduction to Bonding & Chemical Formulas

2. What did the atom of fluorine
say to the atom of
sodium?
You complete me.

3. TERMS
Chemical Bond: force of attraction that holds atoms or ions together
attractive forces determine properties like MPt, Conductivity, and physical state (solid, liquid or gas)
Chemical Compound: two or more elements chemically bonded in the same ratio each time
Chemical Formulas: show the kinds & numbers of atoms in the smallest representative unit of the substance

4. Why do atoms bond?
Changes that  PE and ↑ stability are favored in natural processes.
Compounds are more stable than the uncombined atoms.
gain, lose, and share valence electrons to achieve the octet (OR duet)
gives each atom an electron configuration similar to that of a noble gas
Bond formation releases Energy; spontaneous, exothermic.
Bond breaking absorbs Energy; requires energy input, endothermic.

5. What do I need to know??
The periodic table is a very useful tool!

6. Metals and Nonmetals

7. Types of Chemical Bonding
1. Metal with nonmetal: IONIC
electron transfer = ionic bonding, forms crystals, salts
2. Nonmetal with nonmetal: COVALENT
electron sharing = covalent bonding, forms molecules
Equal sharing= no poles = nonpolar
Unequal sharing = poles form = polar
3. Metal with metal: METALLIC
electron sea , mobile electrons

8. Predicting Bond Type the easy way

9. Ionic Bonds IONIC: e- transfer b/w metals & nonmetals
strong attractive forces b/w ions with opposite charges
bonding elements have large ΔEN
ions bond in an extensive crystal lattice pattern
ions combine to give a neutral compound.
think rocks, salts, minerals

10. Ionic Compounds

11. Using Bohr Models OR Electron Dot Diagrams to represent ionic bonding

12. One representative unit of the crystal lattice = one formula unit = NaCl
Crystal Lattice: the orderly arrangement of ions in an ionic compound.
The chemical formula is always empirical: the simplest combining ratio of the ions

13. Properties of Ionic Substances
Non-conducting, hard, brittle, crystalline solids at room temp.
High melting & boiling points
Will conduct if molten or dissolved in water b/c ions become mobile = electrolytes.

14. Why “Brittle”
A crystal will crack (shatter) when stress (like a hammer) is applied.
Cannot be deformed (not malleable) b/c ions in lattice not free to slide over each other.

15. Why Non-Conducting as a Solid? Why high melting points?
The ions are rigidly trapped in the crystal lattice with very strong forces of attraction between them.
The melting point of NaCl is 801°C!

16. Covalent Bonds = Molecular Compounds
Covalent Bonds are atoms held together by SHARING electrons, in pairs, between NONMETALS, forming molecules.

17. Representing Molecular (Covalent) Compounds
A Lewis structure is a type of structural formula that shows the valence electrons & how they are bonding.

18. Representing Molecular (Covalent) Compounds
A molecular formula shows the types & numbers of atoms in a single molecule.
Chlorine gas: Cl2
Sugar: C6H12O6
Carbon dioxide: CO2
Methane: CH4
Peroxide: H2O2
Ammonia: NH3

19. Covalent Bonds Covalent: e- sharing in pairs b/w two or more nonmetals
bonding atoms have little or no ΔEN
bonding atoms form individual, independent units called molecules
A molecule is a neutral group of atoms united by covalent bonds.
think water, gases, & volatile liquids with vapors

20. Properties of Molecular Compounds
Weak attractive forces between molecules
Liquids or gases at room temperature
Lower Melting Points than Ionic Compounds
Non-conducting in any phase

21. Polarity in molecules
Polar covalent: Uneven charge distribution in a molecule
e- are more strongly attracted to the atom with higher EN.
Results in unequal charge distribution and partial charges, or poles, on the ends of the molecule.
This end has partial positive charge
This end has partial negative charge

22. Polar Covalent Bond
A covalent bond with greater electron density around one of the two atoms H F
electron rich
region
electron poor
region
e- poor
e- rich F H d+ d-

23. Video Clip: Chemical Bonds Cassiopeia Project

24. Seven Diatomic Molecules
A molecule containing only two atoms.
All linear & nonpolar.
Form shape of a “7” on PT
Meet Dr. Brinclhof…. H2 N2 O2 F2
Cl2
Br2 I2

25. Covalent Network Crystals
Covalent bonding between atoms that arrange in a large crystal lattice (instead of individual molecules.)
Ex: Diamonds and Sand (SiO2)
Very hard substances

26. Polyatomic Ions
What looks like a small molecule but acts like an ion???
A polyatomic ion: a charged group of covalently bonded atoms

27. Bond Type by Electronegativity Difference
You can estimate the bond type by subtracting the difference in EN values of the bonding atoms.
Look up EN values & subtract
Hint: how far apart on PT?
∆EN electronegativity difference
Bond
Type
≤0.4
Non Polar Covalent
Between 0.5 to 1.7
Polar Covalent
≥1.9
Ionic

28. Electronegativities (EN)
The ability of an atom in a molecule to attract shared electrons to itself
Linus Pauling

29. Metallic Bonding
All metallic properties can be explained by the sea of flowing electrons characteristic of metallic substances….
Luster: free e- interact with and reflect light
Malleability: can be shaped into sheets without breaking
Conductivity: free e- conduct current

30. Metallic Bonding Occurs within a sample of metal atoms
Results from the attraction between positive metal ions & the surrounding electrons, which are mobile.
The valence e- are delocalized & shared by all the atoms b/c metals have vacant orbitals
This arrangement is not possible in
ionic cmpds: e- are bound in the crystal lattice
covalent cmpds: e- are localized in pairs
Highly mobile sea of electrons accounts for luster, conductivity, malleability and ductility.