1. Chapter 4: Forces Between Particles HCC/TCHS
Chemistry 140
Chapter 4: Forces Between Particles
HCC/TCHS

2. LEARNING OBJECTIVES/ASSESSMENT
When you have completed your study of this chapter, you should be able to:
1. Draw correct Lewis structures for atoms of representative elements.
2. Use electronic configurations to determine the number of electrons gained or lost by atoms as they achieve noble gas configurations.
3. Use the octet rule to correctly predict the ions formed during the formation of ionic compounds, and write correct formulas for binary ionic compounds containing a representative metal and representative nonmetal.
4. Correctly name binary ionic compounds.
5. Determine formula weights for ionic compounds.
6. Draw correct Lewis structures for covalent molecules.
7. Draw correct Lewis structures for polyatomic ions.
8. Use VSEPR theory to predict the shapes of molecules and polyatomic ions .
9. Use electronegativities to classify covalent bonds of molecules, and determine whether covalent molecules are polar or nonpolar.
10. Write correct formulas for ionic compounds containing representative metals and polyatomic ions, and correctly name binary covalent compounds and compounds containing polyatomic ions.
11. Relate melting and boiling points of pure substances to the strength and type of interparticle forces present in the substances.

3. Introduction to Chemical Bonding
Chemical bond – a mutual electrical attraction between the nuclei and valence electrons of different atoms that binds the atoms together
Why are most atoms bonded together?

4. Types of Bonds
Ionic bonding – the electrical attraction between large numbers of anions and cations
Covalent bonding – the result of sharing electron pairs
Metallic bonding – the attraction between metal ions and a sea of electrons
Molecular bonding – bonding between molecules caused by nuclear attraction of outer molecules

5. Naming Binary Compounds
1. First word consists of:
a. a prefix to indicate the number of atoms of the first element
b. the name of the first element
2. Second word consists of:
a. a prefix to indicate the number of atoms of the second element in the formula
b. the root name of the second element
c. the suffix “ide”

6. Nomenclature Prefixes
1 - mono di tri
4 - tetra penta hexa
7 - hepta octa nona
10 - deca

7. Name the following compounds.
3. HBr _______________________
4. CCl4 _______________________
5. As2O5 _______________________
6. N3O4 _______________________
7. Na2O2 _______________________

8. Naming Other Compounds
first word - name of the first species
second word - name of the second species
Roman numerals are used to indicate the oxidation state of transition elements.

9. Name the following compounds.
1. CaCO3 _______________________
2. MgSO4 _______________________
3. H3PO4 _______________________
4. CuNO2 _______________________
5. Fe2(SO4)3 _______________________
6. Fe2(Cr2O7)3 _______________________
7. CuCl _______________________

10. Name the following compounds.
8. FeF2 _______________________
9. NO3 _______________________
10. N2O4 _______________________
11. NH4HSO4 _______________________
12. NaC2H3O2 _______________________
13. H2O _______________________
14. NH3 _______________________

11. Writing Chemical Formulas
Chemical formulas are written from names of compounds.
The sequence for the formula is the same as the name.
The sum of the oxidation numbers of the elements in the formula of a compound must be zero.
polyatomic ion - an ion that consists of two or more atoms acting as a single unit
Parentheses are used to show more than one polyatomic ion.

12. Write formulas for the following.
sodium chloride ____________
magnesium oxide ____________
calcium chloride ____________
ammonium sulfide ____________
aluminum bromide ____________
copper(II) sulfate ____________
potassium peroxide ____________

13. Open Response
Why is it important to have such a detailed system for naming compounds?

14. Why do atoms share electrons?
Hydrogen exists as a molecule and not a single atom? Explain
The proton in one hydrogen atom attracts the electron in the other and vice versa.
What type of force exists between the hydrogen electrons?
The electrons repel each other but this repulsive force is far less than the attractive forces.
covalent bonding - bonding which results from electron sharing

15. Bond Polarity Covalent bonds with uneven electron sharing are polar.
Covalent bonds with even electron sharing are nonpolar.
electronegativity - the tendency for an atom to attract electrons to itself when bonding with other atoms
The difference between electronegativity can be used to indicate bond type.

16. Bond Type To calculate bond type
1. Look up the electronegativity of the atoms in question. (chart passed out)
2. Find the difference between the electronegativities. (subtract)
3. Use the chart to classify the bond as
nonpolar covalent ( )
polar covalent (0.5 – 1.9)
ionic (>2.0)

17. Determine the bond type of each of the following.
NaCl
MgCl2

18. Covalent Bonding and Molecular Compounds
covalent bonding - bonding between atoms which results from electron sharing
The particle which results from covalent bonding is a molecule.
molecule - two or more atoms bonded covalently
There are eight elements in which two atoms bond forming a diatomic molecule. They do not normally exist as single atoms. They are referred to as diatomic elements.
diatomic elements - two identical atoms bonded covalently
molecular compound - a chemical compound whose simplest formula units are molecules

19. Formulas Represent Compounds
chemical formula - a shorthand method of using atomic symbols and subscripts to represent the composition of a substance
molecular formula - a formula indicating the composition of a molecule
dot formula - a formula using dot notation to indicate valence electrons
formula unit – represents the composition of an ionic compound, empirical formula

20. DIATOMIC ELEMENTS NAME DOT FORMULA MOLECULAR FORMULA hydrogen H2
nitrogen N2
oxygen O2
fluorine F2
chlorine Cl2
bromine Br2
iodine I2
astatine At2

21. Lewis Structures
Lewis structures are formulas in which symbols represent nuclei and inner shell electrons and dots, dashes, etc, represent valence electrons.
The Lewis structure of water would be similar to the dot formula.

22. Polyatomic Ions
A polyatomic ion consists of two or more atoms bonded covalently which has a net charge.
Show the Lewis structure for SO4=.
Show the Lewis structure for H2 SO4.

23. Show the Lewis structures for
Examples
Show the Lewis structures for
MgBr2
Na2O
H3PO4
CH3Cl
Al2S3

24. Resonance
Resonance is an attempt to describe bond structure based on data collected about bond length.
Example: ozone, O3

25. Ionic Bonding and Ionic Compounds
Ionic bonding results from electron transfer.
Ion – an atom or group of atoms that has an unbalanced electrostatic charge
Crystal – the particle resulting from ionic bonding
Most ionic compounds are solids.
Salts are examples of ionic compounds.

26. Octet Rule octet - a complete outer shell of eight electrons
Orbital Notation of argon
1s 2s p s p
Ar __ __ __ __ __ __ __ __ __
Show the orbital notation of neon.
1s 2s p
Ne __ __ __ __ __

27. VSEPR Theory
VSEPR - valence shell electron pair repulsion – the valence electron pairs repel each other which moves bonded atoms to an equilibrium position
VSEPR accounts for the bent shape of the water molecule.

28. Molecular Shape
A quick indicator of molecular shape is the number of atoms in a molecule.
The un-bonded electrons must be taken into account to get the exact shape.
What do you think is the shape of: H2
H2O
NH3
CH4

29. Molecular Type
Molecular type is either polar or nonpolar. These are not to be confused with bond type.
Polar molecule – a molecule which lacks symmetry
Nonpolar molecule – a molecule which has symmetry
Which shapes do you expect to be polar and which do you expect to be nonpolar?

30. Rules for Assigning Oxidation Numbers
1. An uncombined element has an oxidation number of zero.
2. A monoatomic ion has an oxidation number equal to its charge.
3. Fluorine has an oxidation number of -1 in all compounds.
4. Oxygen usually has an oxidation number of -2
5. Hydrogen usually has an oxidation number of +1.
6. The more electronegative element in a binary compound is assigned the number equal to the charge it would have if it were an ion.
7. The algebraic sum of the oxidation numbers of the elements in a compound is zero.
8. The algebraic sum of the oxidation numbers of the elements in a polyatomic ion is equal to its charge.

31. Assign oxidation numbers to each element in the following.
1. HCl H ___ Cl___
2. CF4 C ___ F ___
3. PCl3 P ___ Cl___
4. HNO3 H ___ N ___ O ___
5. SiO2 Si___ O ___
6. P2O5 P ___ O ___
7. HClO3 H ___ Cl___ O ___

32. Using Chemical Formulas
The formula mass of sucrose, C12H22O11, is equivalent to the molar mass or the mass of one mole of the compound in grams.
Element # atoms mass total
C x = 144
H x =
O x =
formula mass g/mol

33. Determine the molar mass of each of the following.
1. HNO3
2. Fe2(SO4)3
3. calcium hydroxide
4. barium nitrate
5. CuSO4.5H2O

34. Moles of Compounds
One mole of a compound contains ? molecules or formula units?
The molar mass expressed in grams is the mass of one mole of a compound, Avagadro’s number of formula units or molecules.
Find the mass of 0.25 moles of calcium carbonate.
How many moles in 1Kg of magnesium sulfate?

35. Percentage Composition
Percent can be found by using the formula
part
% = X 100
whole

36. Sample Problem Find the percent of oxygen in water.
1. Write the formula for water.
H2O
2. Find the formula mass of water.
element #atoms mass total H O
-----
18g/mol

37. Sample Problem Continued
3. Calculate % O.
% = (part/whole) X 100
%O = (16/18) X 100
%O = 1600/18
%O = 88.89%
What would be the percent H in water?

38. Find the percentage composition of sodium hydroxide.
1. Write the formula for sodium hydroxide.
2. Find the formula mass.
3. Percentage composition refers to the percent by mass for each element in the formula of the compound. Find the percent Na, O, and H.

39. Calculation of Chemical Formulas from Percentage Composition
1. Assume a 100g sample and express percentages as mass in grams.
2. From the mass data, determine the number of moles of each element present.
3. Write a mole ratio formula.
4. Determine the atom ratio by dividing each mole number by the smallest number and rounding off when appropriate. (Only round 0.1’s and .9’s. If these are not present, multiply through by integers until they are present.)

40. Finding Chemical Formulas
The analysis of a compound determined a percentage composition of 80% carbon and 20% hydrogen. Find the formula.

41. Sample Problems
1. Does smithsonite, ZnCO3, or sphalerite, ZnS, have more zinc per gram of sample?
2. The mineral greenockite is a rare yellow sulfide of cadmium that is 78% cadmium and 22% sulfur by mass. What is the empirical formula of this compound?
3. What is the formula for a compound of aluminum and fluorine that is 32% Al and 68%F?

42. Determination of the Molecular Formula
1. Determine the empirical formula.
2. Determine the empirical formula mass.
3. Use the following formula to determine the multiple, “x”, of the empirical formula.
(empirical formula mass) x = molar mass
4. Multiply the empirical formula subscripts by the multiple “x” to determine the molecular formula.

43. A compound has an empirical formula HO
A compound has an empirical formula HO. The molar mass of the compound is 34g/mol. Find the molecular formula.
1. Empirical formula = HO.
2. Empirical formula mass = = 17g/mol.
3. Determine multiple “x”.
17x = 34
x = 34/17
x = 2
4. Determine molecular formula
(HO)2 = H2O2

44. Sample Problem
The analysis of a compound determined a percentage composition of 80% carbon and 20% hydrogen. The molar mass was determined to be 45 g/mol. Find the molecular formula.
Answer: C3H9