1. CHAPTER SIX (6)
Chemical Bonding

2. Daily Assignment
Explain, using your periodic table, how to calculate (find) the following in an atom:
Atomic number _____________________
Atomic mass _______________________
Identity of the atom _________________
Number of protons __________________
Number of neutrons __________________
Number of electrons ________________

3. The Purpose of an Element
Remember that the desire of every element is to be “happy,” that is, to have a full valence shell.
Atoms are willing to do whatever it takes to accomplish this, including align themselves with some “seedy” atoms
Having a full valence shell is also known as having a stable electron configuration

4. Drawing Elements
In order to talk about bonding, we need to be able to visualize the elements we’re talking about. Here’s how we do that:

5. Ionic Bonding Na gives an electron to Cl, creating two ions. atom atom
When atoms bond together to form compounds, they may gain or lose electrons. An atom that has gained or lost electrons is called an ION.
Na gives an electron to Cl, creating two ions.
atom
atom

6. Let’s Look Closer… The atom Na has 11 (-) electrons, balanced by
11 (+) protons in the nucleus.
The atom Cl has 17 (-) electrons, balanced by 17 (+) protons in its nucleus.

7. THE OCTET RULE
The octet rule says that atoms tend to gain, lose or share electrons so as to have eight electrons in their outer electron shell (HAPPY).
Na atom Cl atom
Na has1 e Cl has 7e-
( it can lose one or gain seven) (it can lose seven or gain one)
If Na gives up its electron to Cl, both will have 8e- .

8. IONS are formed
When a bond is formed, Na gives up one electron to Cl, reducing its number of electrons to 16 and increasing Cl’s electrons to 18.
Na forms a 1+ charge, since it now has one more proton than electron. Cl forms a 1- charge, since it now has one more electron than proton.
The charge is written as a superscript number to the right of the element’s symbol.

9. What is the charge? Mg loses two electrons to form ______________
Oxygen gains two electrons to form __________
Phosphorus gains three electrons to form _______
Carbon gains four electrons to form _________
Calcium loses two electrons to form __________
How was N3- formed? ___________________
How was Ag1+formed? ___________________

10. Daily Assignment
Why do most atoms tend to form compounds, as opposed to existing in their unbonded state?

11. The Structure of an Ion
To find the number of protons, neutrons, and electrons in an ion, you need to look at the atomic number, mass, and the charge on the ion. Remember, atomic number = # of protons atomic mass = protons + neutrons neutrons = atomic mass – atomic number Now, to calculate the number of negative electrons, you must balance the charge of the ions with the number of positive protons.

12. Periodic Table Atomic number = ______ Atomic mass = _______
Charge = ___________
Element name _______
Element symbol _______
# protons __26___
# neutrons __30__
# electrons ___23__
Atomic number = __26__
Atomic mass = _______
Charge = ___________
Element name _______
Element symbol _______
# protons _____
# neutrons __29__
# electrons __24___

13. Periodic Table Atomic number = __26__ Atomic mass = __56___
Charge = ___+3_____
Element name __Iron
Element symbol __Fe___
# protons __26___
# neutrons __30__
# electrons ___23__
Atomic number = __26__
Atomic mass = __55___
Charge = ___+2____
Element name __Iron__
Element symbol __Fe___
# protons __26_
# neutrons __29__
# electrons __24___

14. Periodic Table Atomic number = __16__ Atomic mass = _______
Charge = _____2-____
Element name ______
Element symbol ____
# protons ______
# neutrons __18__
# electrons _____
Atomic number = _____
Atomic mass = ___7____
Charge = ___________
Element name _______
Element symbol _______
# protons __3___
# neutrons _____
# electrons ___2__

15. Periodic Table Atomic number = __16__ Atomic mass = __34__
Charge = _____2-____
Element name ___Sulfur_
Element symbol __S___
# protons __16___
# neutrons __18__
# electrons __18___
Atomic number = __3___
Atomic mass = ___7____
Charge = ____+1____
Element name __Lithium_____
Element symbol __Li___
# protons __3___
# neutrons __4__
# electrons ___2__

16. Periodic Table Atomic number = ______ Atomic mass = __40__
Charge = _____1+____
Element name _______
Element symbol ___K___
# protons _____
# neutrons ____
# electrons _____
Atomic number = __19__
Atomic mass = _______
Charge = ___________
Element name _______
Element symbol _______
# protons _____
# neutrons __22__
# electrons __19___

17. Periodic Table Atomic number = __19__ Atomic mass = __40__
Charge = _____1+____
Element name __Potassium___
Element symbol ___K___
# protons __19__
# neutrons __21_
# electrons __18__
Atomic number = __19__
Atomic mass = ___41____
Charge = _____0______
Element name ___Potassium____
Element symbol ___K___
# protons ___19__
# neutrons __22__
# electrons __19___

18. Review Time!!
Explain the “Octet Rule” _______________________________________ _______________________________________ _______________________________________ ________
One atom has 2 electrons in its outer shell. Another atom has six. Will they combine to form a compound?
What about an atom with two electrons and 2 atoms with seven outer shell electrons?

19. Why do atoms combine?
Atoms will combine with other atoms to achieve a chemically stable arrangement of their electrons.
Recall, the OCTET RULE drives chemical bonding. A chemically stable arrangement of electrons is attainable by completely filling the outer shell (with eight electrons*).

20. To share or to not to share?
Atoms, like small children, will always do what benefits them most In this case, their electrons can either be shared between atoms or transferred from one atom to another to form a chemical bond.

21. Ionic Bond
Ionic Bonds occur when one atom loses electrons to another atom that gains them. Ions arrange themselves into crystal lattices.

22. Ionic Bonds = transfer of electrons between atoms in a reaction
* Only outer shell electrons shown *

23. Writing Ionic Compounds
To write the newly formed ionic compound, take the charges after the reaction, criss-cross them, and drop the sign.
NaCl MgS KI

24. Daily Assignment
Draw a dot diagram and complete the charges for the reaction of:
Sodium Boronate

25. Molecular (Covalent Bond)
Electrons are “shared” between atoms. Since they spend time in the electron clouds of each atom, the atom “feels” complete and happy.

26. The fluorine atoms, being nice to each other as they are, decide to share two electrons - one from each of them. This obviously means that one of the atoms - the one donating - has a complete outer shell; the other atom has a gap for two electrons, e.g. something with six electrons in its outer shell).
Despite the fact that each atom has its own electron cloud, each atom also feels like it has a full outer shell because of the shared electrons.
All electrons are still spinning around each atom, but because of the sharing, the atoms are bonded together in order to maintain the sense of a full outer shell.

27. Molecular (Covalent) Bond
Most atoms become chemically stable by sharing electrons. This does not result in the formation of ions (the atoms do not have charges).
These electrons belong to both H and O. At any point in time they can exist in either electron cloud.

28. Daily Assignment
Do the following reactions at your seat. Draw the dot diagram. Are they covalent or ionic bonds?
Hydrogen and Hydrogen
Chlorine and Chlorine
Hydrogen and Chlorine

29. Daily Assignment
Copy the reactions listed above. Determine the type of chemical bond formed during the reactions.

30. Multiple Covalent Bonds
Sometimes elements do not combine with other elements. Instead, atoms of the same element combine. When this happens, we write it a special way. Take Nitrogen, for example.
Single bond
Double Bond
Triple Bond

31. Unequal Electron Sharing
Those elements on the right of the periodic table crave electrons more than others, and those at the top crave them more than those at the bottom.

32. Polar Covalent Bonds
Occurs when atoms get greedy and “hog” the shared elements.
Example: Water

33. Polar and Nonpolar Molecules
Not all molecules that contain polar bonds are polar molecules. It depends on how the electrons are shared.
Water is a polar molecule because the oxygen has more need for the electrons.
Carbon Dioxide shares the electrons between the carbon and oxygen atoms

34. Polar and Nonpolar Molecules
The attractive forces between polar molecules are greater than those in non polar molecules.
The positive and negative ends of a polar molecule are attracted to eachother and create bonding forces.
Water is an example

35. Writing Chemical Formulas
A chemical formula is a combination of element symbols and subscript numbers that is used to show the composition of a compound.
MgCl2 
The formula above tells you that the compound magnesium chloride contains one atom of magnesium bonded to two atoms of chlorine. There are three total atoms in the compound.

36. Counting Atoms in a Chemical formula
List the name of each element, the number of atoms of that element, and the total number of atoms in the chemical formula.
H2O MgF NH4I

37. Parentheses in a chemical formula
Parentheses are used around polyatomic ions in a chemical formula. The subscript number to the right of the parentheses is used as a multiplier for the atoms inside.
Ba(OH) (NH4)2S Au2(CO3)2

38. Coefficients in a Chemical formula
Coefficients are used in a chemical formula to represent the number of units of the compound present. They act as a multiplier for all atoms present in the formula.
3 H2 O Ba3(PO4)2

39. Counting Atoms in a formulas
List the number of atoms of each element and the total number of atoms in the compound for the formulas listed below.
2 C4H (NH4 )2CO Al2(SO4 )3

40. Daily Assignment
You have learned that elements chemically bond together (ionic or covalent) to form new substances.
You have learned that a chemical formula indicates the number of atoms of each element contained in a chemical compound.
magnesium chloride
MgCl2
Write a formula for a compound with two aluminum and three sulfur atoms. Write a formula for a compound with one beryllium and two iodine atoms.

41. Ionic Compounds calcium fluoride
Ionic compounds are formed from the transfer of electrons, resulting in the creation of a positive ion and a negative ion.
The compound is always written positive ion (cation) then negative ion (anion).
The negative ion (if it is an element) will have an –ide ending.
calcium fluoride

42. Metal Cations
Some metals form multiple ions. Alkalis, Alkaline Earths, and Aluminum form ions with positive charges equal to their group number, such as K+, Ca2+, and Al3+
Many transition metals form more than one type of ion. Examples include copper I and copper II. Roman numerals designate the charge.
Cu2O CuO

43. Polyatomic Ions Ammonium NH4+ Acetate C2H3O2-
Compounds that join together in a regular way and act as one unit. It can have either a negative or positive charge.
Some examples of polyatomic ions:
Ammonium NH4+ Acetate C2H3O2-
Hydroxide OH- Peroxide O22-
Nitrate NO3- Hydrogen Sulfate
Sulfate SO HSO4-
Carbonate CO32- Hydrogen Carbonate
Phosphate PO HCO3-
Chromate CrO42- Hydrogen Phosphate
Silicate SiO32- HPO42-
Permanganate MnO4- Hypochlorate OCl-

44. Writing Compounds Write the symbol for the positive and negative ions.
List the oxidation number (charge) for each. This can be found on your periodic table or listed in the formula name, following the positive ion as a roman numeral.
Find the oxidation numbers for the following ions:
Ag Au (I) Pb (IV) S
Be Pb (II) Al Ca
H Cl O F

45. Criss- cross and drop the charge!
Writing Compounds
Balance the charges by the criss-cross method.
Criss- cross and drop the charge!
Place all parentheses around all polyatomic ions. Make sure the criss-crossed number is placed outside the parentheses.
Reduce subscript numbers to reflect the Least Common Multiple.

46. LETS TRY MORE  Sodium nitride Magnesium bromide Barium sulfide
Lead (IV) oxide
Potassium selenide Chromium nitride Copper (I) oxide Copper (II) oxide

47. Daily Assignment Write the formula for barium chloride.
Explain, using electron dot diagrams, how this compound forms and whether it will be an ionic or covalent bond formed. (I’ll start you off)
Ba + Cl + Cl

48. Writing ionic formulas with polyatomic ions
Place all parentheses around all polyatomic ions. Make sure the criss-crossed number is placed outside the parentheses.
Reduce only the subscript numbers to the outside of the parentheses.
Lead (II) phosphate magnesium acetate
Ammonium sulfide silver silicate

49. Practice with Polyatomic Ions
Aluminum Sulfate Magnesium Acetate
Hydrogen Peroxide Ammonium Nitride
Beryllium Phosphate Calcium Hydrogen Phosphate
Potassium Chromate Potassium Hydroxide

50. More practice with polyatomic ions Take out your homework and periodic table.
Mercury (II) carbonate Platinum (I) sulfide
Platinum (IV) sulfate ammonium nitrate
Calcium iodide hydrogen peroxide

51. Writing formulas of Molecular Compounds
As you recall, elements undergoing molecular (covalent) bonding are sharing electrons, so no charges (or ions) are formed during the formation of the compound.
Molecular compounds are named and written using a system of Greek Prefixes, which must be memorized.

52. Greek Prefixes – Molecular Compounds
One = mono Six = hexa
Two = di Seven = hepta
Three = tri Eight = octa
Four = tetra Nine = nona
Five = penta Ten = deca

53. Writing formulas of Molecular Compounds
Just let the prefix for the element in the formula become the subscript number for that element. If there is no prefix, the subscript is 1.
Examples:
dihydrogen monoxide trisilicon pentabromide
tetraphosphorus hexasulfide carbon dioxide
arsenic trichloride decaphosphorus heptoxide

54. Quiz Chapter Six 1. disulfur hexoxide 2. ammonium sulfide
iron (III) carbonate
iron (II) carbonate
aluminum phosphite
gold (III) nitride
diphosphorus pentabromide
silicon sulfide
tetraradon heptafluoride
potassium oxide

55. Naming Ionic Formulas Write the name of the first element
Add charges to periodic table 
Write the name of the first element
Write the name of the second element changing the name to an _ide ending.
Reverse cris-cross the subscript numbers to determine the charge of both ions.

56. Naming Ionic Formulas
Check the periodic table to determine the number of positive charges that the first element forms. 
If it only forms one, and it is the one listed, then you are done
If it only forms one, but it is not the one listed, then it has been reduced. Determine the multiplier and multiply both charges to determine the correct charge formed.
If it forms more than one, write the Roman numeral equivalent of that charge in the compound’s name.
If it forms more than one, yet none are listed, it has been reduced. See above, and use the Roman numeral.

57. Let’s Try Some 
AlBr3
FeCl3
NiBr2 GeO2

58. Let’s Try Some 
Ni2O3
ZnCl2
PtS2 Cu3P2

59. Naming Ionic Formulas with Polyatomic Ions
Follow all of the same rules, with the following exceptions:
Never change the name of a polyatomic ion
Place parentheses around polyatomic ions before reverse criss- crossing (if they are not already there)
Check the charges of element and polyatomic ions to check for a reduction.

60. Let’s try some 
Au2(CO3)3
FeSO4
Ca3(PO4 )2 (NH4)2CrO4

61. I love naming formulas 
Ba(C2H3O2)2
Pt(NO3)4
Ni2 (SiO3)3 HgSO3

62. Naming Molecular formulas
Simply let the subscript number become the prefix for the element’s name.
The first element keeps its name. It only gets a prefix if it has a subscript in the formula
The second element always gets an – ide suffix (ending). The second element always gets a prefix.
N20 P2O5 S3Cl CO

63. Metallic Bonds
Metals have special bonds because of the way they configure
Metals form regular patterns and share electrons with all the atoms around it, effectively forming cations
This forms a crystalline lattice like those in polar bonds
The metal is effectively neutral, though.

64. Metallic Bonds Properties of Metals explained Conductivity
Metals are good conductors of electricity, but why?
The lattice structure of metal atoms allows electrons to flow freely through the metal
Malleability
A block of ice, if hit with a hammer, would shatter because of its ionic bonds
Although strong, the lattice in a metal is not as rigid as the lattice in ionic bonds. This allows the atoms to move but not break because the electrons keep the metal anions in place

65. Alloys
An alloy is a mixture of metals to give it more desirable properties, such as hardness or resistance to rust
Copper Alloys
First alloy was Bronze – a mixture of copper and tin
Brass – another alloy of copper, formed by mixing copper and zinc

66. Alloys Iron Alloys The best known iron alloy is steel
Steel is a combination of iron and carbon, which hardens the iron.
Stainless steel contains about 10% of chromium, which makes it shiny and resistant to rusting
Other elements, such as phosphorus, manganese, sulfur, and silicon, are added to steel to make strong, flexible metals used in bridges

67. Alloys
Aluminum alloys are used in making airplane sheet metal (mixed with copper or manganese)
Aluminum is also mixed with magnesium to create a lightweight, but strong metal (sometimes used in performance bicycle frames)