1. Bonding and Nomenclature

2. You Must Know Valence Electrons Electron Configuration
Lewis Dot Diagrams

3. Octet Rule Gain or lose electrons to obtain noble gas configuration
It requires energy to gain or lose electrons.
Nature moves toward the LEAST resistence
BECAUSE THIS REQUIRES LESS ENERGY!

4. Oxidation Numbers How do we know? Depends on the nearest noble gas
So the atom will gain or lose electrons
What is easier?
Example: Oxygen has 6 valence electrons
Gain or lose electrons?
Lose 6 or gain 2

5. Survey Says Sulfur Carbon Flourine Hydrogen Potassium Calcium Oxygen
Argon
Carbon
Hydrogen
Calcium
Oxygen
Nitrogen

6. Do you have a Junk drawer at home?
Does it have different kinds of glue?
Paper, Elmer’s, wood, super, etc
Just like you need different types of glue to stick things together, we also need different types of bonds to hold atoms together

7. Chemical Bonding
BOND = Forces of attraction between the nucleus of one atom and the electrons of another atom
ONLY Valence electrons participate in bonding
Bonds are formed as a result of a CHEMICAL REACTION

8. SO why is chemical bonding so important?
#1 Used when creating chemical compounds Important because most things are made of these compounds #2 Determine different physical properties of compounds they create

9. SO why is chemical bonding so important?
#3 Also determine chemical properties #4 Chemical compounds that are formed by chemical bonds help make the body function as it does Ex: stomach would not be able to digest food

10. SO why is chemical bonding so important?
#5 With the use of chemical bonds, we create chemical compounds that allow us to find new substances. These substances can be used in new tech, medicine, and products. Chemical bonds are what hold the world together

11. Energy & bonding Energy associated with bonding:
Exothermic (spontaneous)
&
Endothermic (not spontaneous)

12. Exothermic Spontaneous bond formation = energy is released
go from HIGH energy (unhappy atoms) to Lower energy (happy atoms)
Creating a bond creates STABILITY
Energy is RELEASED as a product
A + B → C + energy

13. Endothermic Breaking Bonds (not spontaneous) = Energy is consumed
Go from LOW energy (happy atoms) to HGHER energy (unhappy atoms)
Ripping two atoms apart takes energy
Energy is CONSUMED or needed as an ingredient to fuel the process
A + energy → B C

14. Keeping Track of Electrons
The electrons responsible for the chemical properties of atoms are those in the outer energy level.
Valence electrons - The s and p electrons in the outer energy level.
Core electrons -those in the energy levels below.

15. Keeping Track of Electrons
Atoms in the same column
Have the same outer electron configuration.
Have the same valence electrons.
Easily found by looking up the group number on the periodic table.
Group 2A - Be, Mg, Ca, etc.-
2 valence electrons

16. X Electron Dot diagrams A way of keeping track of valence electrons.
How to write them
Write the symbol.
Put one dot for each valence electron
Don’t pair up until they have to X

17. The Electron Dot diagram
for Nitrogen
Nitrogen has 5 valence electrons.
First we write the symbol. N
Then add 1 electron at a time to each side.
Until they are forced to pair up.

18. Write the electron dot diagram forNa
1s22s22p63s1 Na Mg C O F Ne He Mg
1s22s22p63s2 C
1s22s22p2 O
1s22s22p4 F
1s22s22p5 Ne
1s22s22p6 He
1s2

19. Electron Configurations for Cations
Metals lose electrons to attain noble gas configuration.
They make positive ions.
If we look at electron configuration it makes sense.
Na 1s22s22p63s1 - 1 valence electron
Na+ 1s22s22p6 -noble gas configuration

20. Electron Dots For Cations
Metals will have few valence electrons
These will come off
Forming positive ions Ca
40.078 20
1s22s22p63s2 Ca +2

21. Electron Configurations for Anions
Nonmetals gain electrons to attain noble gas configuration.
They make negative ions.
If we look at electron configuration it makes sense.
S 1s22s22p63s23p4 - 6 valence electrons
S-2 1s22s22p63s23p6 -noble gas
configuration.
(anions)

22. Electron Dots For Anions
Nonmetals will have many valence electrons.
They will gain electrons to fill outer shell. P
P-3

23. Stable Electron Configurations
All atoms react to achieve noble gas configuration.
Noble gases have two s and six p electrons.
Eight valence electrons .
Also called the octet rule. Ar

24. Four Types of Naming Binary compounds Ternary compounds
Coordination compounds
Organic compounds
Contain only two types of elements
Contain more than two types of elements
These will not be covered
We will cover these in a separate unit

25. Types of Bonding
Ionic & Covalent

26. Guiding Questions? What is that?
How do we figure out what the chemical formula is?
What does it mean to be "free of chemicals"?
Chemical formulas can be predicted from the periodic table and allow chemists to classify and predict properties of compounds
The general trend in the universe to strive for lower energy explains and allows for prediction of chemical properties of elements
Elements combine in whole number ratios and these molar ratios can be used to determine chemical formulas
More Specifically...:
Bonding Model
Use physical and chemical properties to distinguish between ionic and covalent compounds
Describe energy changes as elements combine to form an ionic compound
Describe ionic bonding as the transfer of electrons and the formation of a crystal lattice due to electrostatic attraction between ions of opposite charge
Predict the formation of cations and anions based on placement on periodic table
Relate formation of anion or cation with ionization energy and electron affinity
State that bonding occurs to increase stability
Contrast metallic and ionic bonding
Nomenclature and formulas
Learn the names and formulae of common anions and cations, including carbonate, sulfate, nitrate, hydroxide, phosphate, and ammonium
Write chemical formulas for ionic compounds given
a. Name of compound or
b. A pair of ions
Identify polyatomic ions
Classify compounds as being ionic or covalent.
Name ionic compounds using stock system (Roman numerals)
Math
Calculate molecular mass
Use dimension analysis to convert between moles, grams, atoms, ions and molecules

27. General Rule of Thumb Metal + nonmetal = Ionic Bond
Polyatomic ion + metal or polyatomic ion = Ionic bond
Nonmetal + Nonmetal = Covalent Bond

28. Ionic Bonding
Formed from: Metal & Nonmetal Name cation first followed by anion

29. Ionic Bonding
Anions and cations are held together by opposite charges. (Neutral)
Ionic compounds are called salts.
Simplest ratio is called the formula unit.
The bond is formed through the transfer of electrons.
Electrons are transferred to achieve noble gas configuration.

30. Properties of Ionic Compounds
Crystalline structure.
A regular repeating arrangement of ions in the solid.
Ions are strongly bonded.
Structure is rigid.
High melting points- because of strong forces between ions.

31. Ionic Bonding
Compounds formed by the TRANSFER of electrons from one atom to another
Metal loses e− to a Nonmetal 𝑁𝑎 𝐶𝑙 −
Metal combines with a Polyatomic ion 𝑁𝑎 𝑁𝑂 3
Polyatomic ion combines with a Nonmetal 𝑁𝐻 𝐶𝑙 −
Polyatomic ion combines with another Polyatomic ion
𝑁𝐻 𝑁𝑂 3

32. Formula writing for Ionic Compounds
IUPAC = International Union of Pure and Applied Chemistry
Compounds have a common name and a chemical name
There is a systematic method for naming ionic compounds
Need 2 types of ions and the CHARGE of each ion.

33. Criss-cross rule Name the Cation first (Metal
Name the anion second (Nonmetal)
Replace the ending of the anion with “IDE”
Ex: Flourine – Flouride
oxygen - oxide

34. Example: Aluminum Chloride
Criss-Cross Rule
Step 1:
Aluminum Chloride
write out name with space
Al Cl 3+ 1-
Step 2:
write symbols & charge of elements
Al Cl
Step 3: 1 3
criss-cross charges as subsrcipts
Step 4: AlCl 3
combine as formula unit
(“1” is never shown)

35. Example: Aluminum Oxide
Criss-Cross Rule
Step 1: Aluminum Oxide
Step 2: Al3+ O2-
Step 3: Al O 2 3
Step 4: Al2O3

36. Ionic Bonding
transfer of electron + - Na Cl
NaCl

37. Ca +2 Ca P -3 +2 Ca P -3 +2 Ionic Bonding
All the electrons must be accounted for!

38. Ionic Bonding
Bonding Activity

39. BrO41- BrO31- BrO21- BrO1- CO42- CO32- CO22- CO2- ClO41- ClO31- ClO21-
Polyatomic Ion:
a group of atoms that stay together and have a single, overall charge.
BrO41-
Perbromate ion
BrO31-
Bromate ion
BrO21-
Bromite ion
BrO1-
Hypobromite ion
CO42-
CO32-
Carbonate ion
CO22-
CO2-
ClO41-
ClO31-
Chlorate ion
ClO21-
ClO1-
IO41-
IO31-
Iodate ion
IO21-
IO1-
NO41-
NO31-
Nitrate ion
NO21-
NO1-
PO53-
PO43-
Phosphate ion
PO33-
PO23-
SO52-
SO42-
Sulfate ion
SO32-
SO22-
1 more oxygen
“normal”
1 less oxygen
2 less oxygen

40. Polyatomic Ions - Memorize
Eight “-ATE’s”
PO43- ……………
SO42- ……………
CO32- …………..
ClO31- …………..
NO31- ………..….
phosphate
sulfate
carbonate
chlorate
nitrate
phosphATE
sulfATE
carbonATE
chlorATE
nitrATE
Exceptions:
ammonium
hydroxide
cyanide
NH41+ ……………
OH1- ……………
CN1- …………..

41. Formula writing for Ionic Compounds
Polyatomic ions are named exactly as they are seen on Table
Parentheses are REQUIRED only when you need more than one “bunch” of a particular Polyatomic ion.
Examples:
NaOH – sodium hydroxide KNO3 – Potassium nitrate Ammonium hydroxide – NH4OH Calcium Phosphate – CaP
𝑀𝑔 +2 and 𝑁𝑂2 − - Mg(NO2) 𝑁𝐻 4 and CLO (NH4)3 N

42. Ionic bonding
Transition metals tend to have more than one oxidation state.
Use roman numerals to indicate their oxidation #.
Roman numeral appears in ( ) AFTER the element symbol

43. Covalent bonding Formed from 2 or more Nonmetals
Atoms SHARE e to get a full valence shell
GOAL… to have a full outer shell
Exception: H needs 2 𝑒 −
Examples: 𝐶𝐻 & 𝐻 2 O

44. Oxidation States in Formulas and Names
Traditional System Stock System
(Two non-metals)
dinitrogen monoxide N2O nitrogen (I) oxide
dinitrogen trioxide N2O nitrogen (III) oxide
dinitrogen pentoxide N2O5 nitrogen (V) oxide
sulfur dioxide SO sulfur (IV) oxide
sulfur trioxide SO sulfur (VI) oxide
stock system is NOT preferred for two non-metals

45. F F Covalent bonding A second atom also has seven By sharing electrons
Fluorine has seven valence electrons
A second atom also has seven
By sharing electrons
…both end with full orbitals
8 Valence
electrons
8 Valence
electrons F F

46. Single Covalent Bond A sharing of two valence electrons.
Only nonmetals and Hydrogen.
Different from an ionic bond because they actually form molecules.
Two specific atoms are joined.
In an ionic solid you can’t tell which atom the electrons moved from or to.

47. Covalent compounds FORGET Charges
Use greek prefixes to indicate how many atoms of each element, but do not use mono on first element.
Examples:
Carbon dioxide CO2
Dinitrogen trioxide - N2O3
1 – mono
6 – hexa
2- di
7 – hepta
3 – tri
8 – octa
4 – tetra
9 – nona
5 – penta
10 - deca

48. Rules for writing Covalent compounds
Least electronegative element is written first
Most electronegative element is written last
Subscripts tell you the prefix of each element in the formula

49. How to show how they formed
It’s like a jigsaw puzzle.
I have to tell you what the final formula is.
You put the pieces together to end up with the right formula.
For example- show how water is formed with covalent bonds.

50. Water H
Each hydrogen has 1 valence electron Each hydrogen wants 1 more The oxygen has 6 valence electrons The oxygen wants 2 more They share to make each other happy O

51. H O Water Put the pieces together The first hydrogen is happy
The oxygen still wants one more H O

52. H O H O H Water The second hydrogen attaches
Every atom has full energy levels
A pair of electrons is a single bond H O H O H

53. Lewis Structures 1) Count up total number of valence electrons
2) Connect all atoms with single bonds
- “multiple” atoms usually on outside
- “single” atoms usually in center;
C always in center,
H always on outside.
3) Complete octets on exterior atoms
(not H, though)
4) Check
- valence electrons math with Step 1
- all atoms (except H) have an octet;
if not, try multiple bonds
- any extra electrons?
Put on central atom

54. How to draw them Add up all the valence electrons.
Count up the total number of electrons to make all atoms happy.
Subtract.
Divide by 2
Tells you how many bonds - draw them.
Fill in the rest of the valence electrons to fill atoms up.

55. Multiple Bonds
Sometimes atoms share more than one pair of valence electrons.
A double bond is when atoms share two pair (4) of electrons.
A triple bond is when atoms share three pair (6) of electrons.

56. C O Carbon dioxide CO2 - Carbon is central atom ( I have to tell you)
Carbon has 4 valence electrons
Wants 4 more
Oxygen has 6 valence electrons
Wants 2 more C O

57. Carbon dioxide
Attaching 1 oxygen leaves the oxygen 1 short and the carbon 3 short C O

58. Carbon dioxide
Attaching the second oxygen leaves both oxygen 1 short and the carbon 2 short O C O

59. O C O Carbon dioxide The only solution is to share more
Requires two double bonds
Each atom gets to count all the atoms in the bond
8 valence electrons
8 valence electrons
8 valence electrons O C O

60. N H Examples NH3 N - has 5 valence electrons wants 8
H - has 1 valence electrons wants 2
NH3 has 5+3(1) = 8
NH3 wants 8+3(2) = 14
(14-8)/2= 3 bonds
4 atoms with 3 bonds H

61. H H N H Examples Draw in the bonds All 8 electrons are accounted for
Everything is full H H N H

62. Examples HCN C is central atom N - has 5 valence electrons wants 8
C - has 4 valence electrons wants 8
H - has 1 valence electrons wants 2
HCN has = 10
HCN wants = 18
(18-10)/2= 4 bonds
3 atoms with 4 bonds -will require multiple bonds - not to H

63. H C N HCN Put in single bonds Need 2 more bonds
Must go between C and N H C N

64. H C N HCN Put in single bonds Need 2 more bonds
Must go between C and N
Uses 8 electrons - 2 more to add H C N

65. H C N HCN Put in single bonds Need 2 more bonds
Must go between C and N
Uses 8 electrons - 2 more to add
Must go on N to fill octet H C N

66. Another way of indicating bonds
Often use a line to indicate a bond
Called a structural formula
Each line is 2 valence electrons H O H H O H =

67. H C N H C O H Structural Examples
C has 8 electrons because each line is 2 electrons
Ditto for N
Ditto for C here
Ditto for O H
C O H

68. Resonance
When more than one dot diagram with the same connections are possible.
NO2-
Which one is it?
Does it go back and forth.
It is a mixture of both, like a mule.
NO3-

69. Electronegativity
A measure of how strongly the atoms attract electrons in a bond.
The bigger the electronegativity difference the more polar the bond.
Covalent nonpolar
Covalent moderately polar
Covalent polar
> 2.0 Ionic

70. POLAR VS NONPOLAR Polar Nonpolar Asymmetrical molecules
UNUEQUAL sharing of electrons
Does not pass the “mirror test”: can not be folded to reflect itself
EQUAL sharing of electrons
Does pass the “mirror test”: can be folded to reflect itself
2 atoms different Elements (EN)
2 atoms same element (EN)
> 2 atoms unbonded e or lone pairs around central atom
> 2 atoms No unbonded e or lone pairs around the central atom
Ex: HCl 𝐻 2 O
Ex: 𝐶𝑙 𝐶𝑂 𝐶𝐶𝑙 4
BEWARE! There are often POLAR bonds inside NONPOLAR molecules

71. Polar Bonds
When the atoms in a bond are the same, the electrons are shared equally.
This is a nonpolar covalent bond.
When two different atoms are connected, the atoms may not be shared equally.
This is a polar covalent bond.
How do we measure how strong the atoms pull on electrons?

72. How to show a bond is polar
Isn’t a whole charge just a partial charge
d+ means a partially positive
d- means a partially negative
The Cl pulls harder on the electrons
The electrons spend more time near the Cl d+ d- H Cl

73. Polar Molecules
Molecules with ends

74. Must determine geometry first.
Polar Molecules
Molecules with a positive and a negative end
Requires two things to be true
The molecule must contain polar bonds
This can be determined from differences in electronegativity.
Symmetry can not cancel out the effects of the polar bonds.
Must determine geometry first.

75. Is it polar? HF
H2O
NH3
CCl4
CO2

76. Intermolecular Forces
What holds molecules to each other

77. Intermolecular forces
ONLY in covalent molecules
Weak forces that act between molecules that hold molecules to each other
Only exist in gaseous & liquid states
Called weak forces because they are much weaker than chemical bonds
REMEMBER: IMF’s occur b/t molecules, whereas
BONDING occurs within molecules

78. Other types of bonds London dispersion forces (LDF’s)
Dipole (dipole-dipole)
Hydrogen Bonds

79. Intermolecular Forces
They are what make solid and liquid molecular compounds possible.
The weakest are called van derWaal’s forces - there are two kinds
Dispersion forces
Dipole Interactions
depend on the number of electrons
more electrons stronger forces
Bigger molecules

80. London dispersion Weakest of all the IMF’s
Only important for NONPOLAR molecules
More electrons = Greater LDF’s
Electron-electron repulsion creates brief dipoles in atoms and molecules.

81. Dipole interactions
Occur when polar molecules are attracted to each other.
Slightly stronger than dispersion forces.
Opposites attract but not completely hooked like in ionic solids.

82. Dipole interactions Fluorine is a gas Bromine is a liquid
Depend on the number of electrons
More electrons stronger forces
Bigger molecules more electrons
Fluorine is a gas
Bromine is a liquid
Iodine is a solid

83. H F d+ d- H F d+ d- Dipole interactions
Occur when polar molecules are attracted to each other.
Slightly stronger than dispersion forces.
Opposites attract but not completely hooked like in ionic solids.
H F
d+ d-
H F
d+ d-

84. Dipole-Dipole
Molecules such as HCl have both POSITIVE and a NEGATIVE end or POLES
Two poles = DIPOLE
Results from an UNEQUAL/ASYMMETRICAL sharing of electrons
Dipole-Dipole = 2 molecules with permanent dipoles are attracted to one another temporarily.

85. Dipole-dipole
Dipole moment = measure of the strength of the dipole within a molecule (Polarity)
The greater the difference in electronegativity between atoms, the greater the polarity/dipole moment.
The higher the dipole moment, the stronger the intermolecular forces.
The stronger the IMF’s the higher the melting and boiling point.

86. Dipole Interactions
d+ d-
d+ d-
d+ d-
d+ d-
d+ d-
d+ d-
d+ d-
d+ d-

87. Hydrogen bonding Specific type of DIPOLE interaction
In a Polar bond, hydrogen is basicallly reduced to a bare proton with almost no atomic radius.
Strongest of all IMF’s by far
ONLY occur in molecules containing hydrogen and fluorine, Oxygen, or Nitrogen.

88. Hydrogen bonding
Are the attractive force caused by hydrogen bonded to F, O, or N.
F, O, and N are very electronegative so it is a very strong dipole.
The hydrogen partially share with the lone pair in the molecule next to it.
The strongest of the intermolecular forces.

89. Hydrogen Bonding H O d+ d- H O d+ d-

90. Hydrogen bonding H O H O H O H O H O H O H O

91. Metallic Bonds
In metals, valence shells of atoms overlap and are free to travel between atoms through material.
Atoms of a metal DO NOT bond with other metal atoms.
Metals share a sea of MOBILE valence electrons
This allows metals to conduct electric current.