1. Chapter 5-The Structure of Matter
Molecules and Compounds
Ionic and Covalent Bonding
Empirical and Molecular Formula

2. Compounds and Molecules

3. Remember from Chapter 2…
A compound is a substance made of atoms of two or more different elements that are chemically combined by chemical bonds
Mixtures are made of different substance that are just placed together; each substance in the mixture keeps its own properties

4. What is a chemical bond?
An attractive force between the nuclei and valence electrons of different atoms that binds them together
Bonds form in order to…
decrease potential energy (PE)
increase stability

5. Ne Stability Stability is the driving force behind bond formation!
Octet Rule
most atoms form bonds in order to have 8 valence e-
full outer energy level
like the Noble Gases! Ne
Stability is the driving force behind bond formation!

6. Stability
Transferring e-
Sharing e-

7. H2O Chemical Formulas 1 oxygen atom 2 hydrogen atoms
A compound always has the same chemical formula
Shows:
Elements in the compound
The ratios of elements in the compound
Example:
H2O
1 oxygen atom
2 hydrogen atoms

8. Chemical Structure This tells you how the atoms or ions are arranged
Described by bond length and bond angles
Bond length: distance between the nuclei of two bonded atoms
Bond angles: tell you how far away the atoms are away from each other in relation to the center atom
Bond angles try to be as large as possible to decrease stress
CO2 has bond angles of 180 degrees
BF3 has bond angles of 120 degrees
CH4 has bond angles of degrees

9. Models of Compounds Space-filling: Ball and Stick:
Shows the relative space occupied by oxygen and hydrogen
Ball and Stick:
can give you an idea of the geometry of a compound

10. How does structure affect properties?
Some compounds exist as a large network of bonded atoms like quartz
Rigid, strong structures
Have high boiling and melting points
Some compounds exist as a large network of bonded positive and negative ions
Crystals are cubed shaped
In NaCl, made of repeated network connected by strong bonds
Made of tightly packed + and – ions
Strong attractions cause high melting and boiling points
Some compounds are made of many separate molecules
Less strongly attracted atoms
The strength of these attraction varies greatly

11. NaCl CO2 Vocabulary IONIC COVALENT Formula Unit Molecular Formula
CHEMICAL FORMULA
IONIC
COVALENT
Formula
Unit
Molecular
Formula
NaCl
CO2

12. NaCl NaNO3 A. Vocabulary more than 2 elements 2 elements Binary
COMPOUND
more than 2
elements
2 elements
Binary
Compound
Ternary
Compound
NaCl
NaNO3

13. Na+ NO3- A. Vocabulary 1 atom 2 or more atoms Monatomic Ion Polyatomic

14. Types of Bonds

15. Ionic Bonds Ionic Bonds are formed by the TRANSFER of electrons
What constitutes an ion?
Positive charge=cation
Negative charge=anion
Opposites attract
Cations will attract anions
Anions will attract cations
Bonding ALWAYS takes place in the valence shell between valence electrons

16. Properties of ionic bonds
Strong bonding forces
Hard, brittle
Crystalline structure
Conduct electricity in solution
High melting/boiling point

17. What makes an ionic bond?
Must include a metal + a non-metal
Metal = cation
Non metal = anion
Anion “takes” electron from cation
Cation is willing to give it away, to lower an energy level
Anion is looking to fulfill the octet rule
For this reason – group I and II metals “LOVE” group 17 non-metals (halogens)
Creates a salt

18. Covalent Bonds Compounds that are made of molecules
Bonds that are formed by SHARING electrons
Electrons are shared from the valence shell, still want it full!
But electrons are not always shared equally…
When two Chlorine atoms bond, electrons are equally attracted to the positive nucleus of each atom---NON- POLAR Covalent bond
When not shared equally-Polar Covalent bond

19. Properties of Covalent Bonds
Strong bonding forces
Can be found as liquids, gases, or soilds
Do not conduct electricity
Low melting/boiling point
Non-metal + non-metal
All the polyatomics are covalently bonded

20. What makes a Covalent Bond
Non-metal + non-metal
Some examples below
Can share more than one pair of electrons
single bond shares one pair electrons
double bond shares two pairs electrons
triple bond shares three pairs electrons

21. Naming Covalent Compounds
Write the more metallic element first.
Add subscripts according to prefixes.

22. Covalent Bonds
Two nonmetal atoms form a covalent bond because they have less energy after they bonded
H H H : H = HH = H2
hydrogen molecule

23. Diatomic Molecules Gases that exist as diatomic molecules
Are: Br2, I2 , N2, Cl2, H2, O2, F2
octets
       
 N   N   N:::N
 
triple bond

24. Molecular Formulas
The Seven Diatomic Elements
Br2 I2 N2 Cl2 H2 O2 F2

25. Comparison Chart IONIC COVALENT Bond Formation The “glue” of the bond
e- are shared between two nonmetals
e- are transferred from metal to nonmetal
The “glue” of the bond
electrical attraction
Physical
State
solid
liquid , gas, or solid
Melting
Point
high
low
Solubility in
Water
yes
usually not
Electrical Conductivity
yes no
Other
Properties
crystal lattice shape = brittle; exothermic
odor

27. Metallic Bonds Metal + metal
Attraction between one atom’s nucleus and a neighboring atom’s electrons
Closely packed structure, causes outermost energy level (valence shell) to overlap
Electrons are free to move from atom to atom
This is why they are good conductors of electricity

28. Metallic Bond Chart METALLIC “electron sea” solid very high no
e- are delocalized among metal atoms
Bond Formation
“electron sea”
Type of Structure
Physical State
solid
Melting Point
very high
Solubility in Water no
Electrical Conductivity
yes (any form)
Other
Properties
malleable, ductile, lustrous

29. Hydrogen Bonds
Explains why an oxygen of one water molecule is attracted to hydrogen atom of a neighboring water molecule
Water molecules attract each other, but these bonds aren’t as strong as the bonds holding oxygen and hydrogen molecules together within a molecule

30. Writing Lewis Dot Structures for Covalent Compounds

31. Lewis dot structres Na Cl =
Simple depiction to help us “visualize” the location of the valence electrons
Only includes valence electrons
Includes element symbol and dots representing valence electrons Na Cl =
Loses electron; becomes +
Gains electron; becomes -

32. What’s the charge? The charge on an ion.
Indicates the # of e- gained/lost to become stable. 1+ 4- 2+ 3+ 4+ 3- 2- 1-

33. How do ions bond? +1 -1 Review: cations + anions
What is the ionic charge of Na? +1
What is the ionic charge of Cl? -1
The final compound MUST be neutral!!

34. NASL Method
STEPS
Figure out what element is central (I will tell you this part)
Calculate the N (needed) number of electrons for all atoms to abide by the octet rule (except 2 for H & He)
Calculate the A (available) number of electrons by adding up all the valence electrons for each atom
Calculate the S (shared) number of electrons = N-A
Calculate lone L (lone pairs or dots) as the difference between A – S.
Check to make sure you have used as many electrons in A step

35. H2O N A S L

36. SO2 N A S L

37. SOCl2 N A S L

38. O2 N A S L

39. What if it’s charged? STEPS When Calculating N include charge
Anions GAIN that many electrons
Cations LOSE that many electrons
Calculate A, S, L just the same
When finished with lewis dot structures, draw brackets around picture with the charge on the outside

40. (NH4)+1 N A S L

41. (NO3)-1 N A S L

42. (SO4)-2 N A S L

43. Empirical and Molecular Formulas

44. C2H6 CH3 Empirical Formula
Smallest whole number ratio of atoms in a compound
C2H6
reduce subscripts
CH3

45. Empirical Formula 1. Find mass (or %) of each element.
2. Find moles of each element.
3. Divide moles by the smallest # to find subscripts.
4. When necessary, multiply subscripts by 2, 3, or 4 to get whole #’s.

46. B. Empirical Formula
Find the empirical formula for a sample of 25.9% N and 74.1% O.
25.9 g
1 mol
14.01 g
= 1.85 mol N
= 1 N
1.85 mol
74.1 g
1 mol
16.00 g
= 4.63 mol O
= 2.5 O

47. N2O5 N1O2.5 B. Empirical Formula
Need to make the subscripts whole numbers
 multiply by 2
N2O5

48. CH3 C2H6 C. Molecular Formula empirical formula molecular formula ?
“True Formula” - the actual number of atoms in a compound
Different compounds can have the same empirical formula
CH3
empirical
formula ?
C2H6
molecular
formula

49. C. Molecular Formula 1. Find the empirical formula.
2. Find the empirical formula mass.
3. Divide the molecular mass by the empirical mass.
4. Multiply each subscript by the answer from step 3.

50. (CH2)2  C2H4 C. Molecular Formula empirical mass = 14.03 g/mol
The empirical formula for ethylene is CH2. Find the molecular formula if the molecular mass is g/mol?
empirical mass = g/mol
28.1 g/mol
14.03 g/mol
= 2.00
(CH2)2  C2H4