1. Chapters 7-9: Chemical Composition

2. Intra-chemical Forces
Intra = within
Atoms (elements) held together by an
attractive force

3. Metallic Bonding
atoms of a metal “share” the valence electrons because they move from one element to another

4. Ionic Bonding valence electrons are transferred between two elements
strongest bonds

5. Covalent Bonding valence electrons are shared between two elements
Weaker bonds than ionic bonds

6. Ionic Bonding
Electrons are transferred from one element to another.

7. Ionic Bonding Opposite charges is attractive force
Commonly referred to as “salts”
Atoms that donates electron = cation
Atom that accepts electron = anion
Oxidation state refers to the charge of an atom

8. Lewis Dot Formulas
Octet Rule: every element wants 8 electrons in its outer shell.
potassium + chlorine → potassium chloride
magnesium + fluorine → magnesium fluoride

9. Types of Ions monatomic cation: cation with one element
K+ Mg Fe Fe Mn Au Au+
monatomic anion: anion with one element
name ends in – ide
Cl– O2– N3– S2– F– P3– Br –
polyatomic ion: many atoms covalently bonded that have a net charge.
NO3– SO42– C2H3O2– PO43– NH4+

10. Writing Ionic Chemical Formulas
Composition
number of elements
Writing chemical formulas (from the names)
Recognize the (+) and (–) ions
Write the symbols of the elements with their charge
A Roman numeral will tell you what the charge is on the cation if there is more than one possibility
Adjust the number of each ion (with subscripts) as needed so the positive charge is equal and opposite the negative charge.
If the ions are polyatomic and there is more than one, the ion is enclosed with parentheses with a subscript on the outside.

11. Writing Ionic Chemical Formulas
sodium chloride
calcium sulfide
calcium sulfate
barium phosphate

12. Naming Ionic Compounds
Consists of two words:
Name the cation
Name the anion
If the cation has more than one possible charge, a Roman Numeral is used to show the charge.
All transition metals need roman numerals except:
Zinc always has a charge of +2
Silver always has a charge of +1

13. Naming Ionic Compounds
FeCl3 Fe3+  iron(III) chloride
FeCl2 Fe2+  iron(II) chloride
NH4Cl
Cu2SO4
NaC2H3O2
Ca(NO3)2
Zn(ClO)2
Cu2O
CuO

14. Naming Ionic Compounds
FeCl3 Fe3+  iron(III) chloride
FeCl2 Fe2+  iron(II) chloride
NH4Cl ammonium chloride
Cu2SO4 copper (I) sulfate
NaC2H3O2 sodium acetate
Ca(NO3)2 calcium nitrate
Zn(ClO)2 zinc hypochlorite
Cu2O copper (I) oxide
CuO copper (II) oxide

15. Covalent Bonding
Valence electrons are shared between two elements
Weaker than ionic bonding

16. polar & nonpolar covalent bonds
Polar Covalent (stronger): unequal sharing of electrons (the more electronegative element pulls more)
Nonpolar Covalent (weaker): equal sharing of electrons

17. Writing Formulas for Covalent Compounds
carbon dioxide
carbon monoxide
dinitrogen monoxide
carbon tetrafluoride
triphosphorus pentachloride

18. Naming Formulas for Covalent Compounds
Binary covalent compounds (2 elements)
Formulas with two nonmetals
Rules:
First word:
prefix indicating the number of atoms for the first element (if there is more than one)
name of first element
Second word:
prefix for the number of atoms of the second element (prefixes on supplement notes sheet)
name of second element
suffix –ide

19. Naming Formulas for Covalent CompoundsNO
NO2
CBr4
P4O10
BF3
SiI5
H2O
S6Cl8
Se7O9

20. Naming Formulas for Covalent Compounds
NO nitrogen monoxide
NO2 nitrogen dioxide
CBr4 carbon tetrabromide
P4O10 tetraphosphorus decoxide
BF3 boron trifluoride
SiI5 silicon pentaiodide
H2O dihydrogen monoxide
S6Cl8 hexasulfur octochloride
Se7O9 heptaselenium nonoxide

21. Lewis Structures
The number of covalent bonds formed by an atom equals the number of unpaired electrons in the Lewis Dot Formula.
i. water (H2O)

22. Lewis Structures
ii. Hydrogen gas (H2)
iii. Hydrochloric acid (HCl)

23. Lewis Structures
iv. ammonia (NH3)
v. methane (CH4)

24. i. double bonds: two pairs of electrons shared O2
Multiple Bonds
i. double bonds: two pairs of electrons shared O2
ii. triple bonds: three pairs of electrons shared N2

25. Hybridization
Combining of two or more orbitals of nearly the same energy into new orbitals of equal energy

26. Beryllium: [He]2s2 sp hybrid
Hybridization
Most common hybridizations occur in groups 2, 13, 14 (IIA, IIIA, IVA)
Group 2 (IIA):
Beryllium: [He]2s2 sp hybrid

27. Boron: [He]2s22p1 sp2 hybrid
Hybridization
Most common hybridizations occur in groups 2, 13, 14 (IIA, IIIA, IVA)
Group 13 (IIIA):
Boron: [He]2s22p1 sp2 hybrid

28. Carbon: [He]2s22p2 sp3 hybrid
Hybridization
Most common hybridizations occur in groups 2, 13, 14 (IIA, IIIA, IVA)
Group 14 (IVA):
Carbon: [He]2s22p2 sp3 hybrid

29. Molecular Polarity
Molecules with more than one element (polar or nonpolar)
depends on:
i. electronegativity difference (2 elements)
ii. Non-bonded electron pairs (2+ elements)
iii. Structure (symmetry) (2+ elements)

30. “Inter-chemical” forces
A. Inter = between
B. Whole salts or molecules attract and bond with one another

31. “Inter-chemical” forces
Ion – dipole
Dipole – Dipole
Hydrogen Bonding
London Dispersion

32. Ion – Dipole forces
Strongest inter-chemical force

33. hydrogen bonding is a unique case of dipole – dipole bonding
occurs because hydrogen’s exposed proton results in a slight positive charge.

34. hydrogen bonding medium strength inter-chemical bond.
occurs in molecules when hydrogen is bonded with F, O, or N.

35. hydrogen bonding hydrogen bonding is responsible for:
water’s high boiling point,
and the low density of ice

36. dipole – dipole bonding
weaker than hydrogen bonding.
occurs between polar molecules

37. London Dispersion Forces
named after Fritz London
the weakest
inter–molecular force
the random movement of electrons can create an instantaneous dipoles