1. Bonding Unit -Learn to draw Lewis Structures.
-Define Ionic, Polar and Nonpolar Covalent Bonding, metallic bonding.
-Discuss ionic, covalent and metallic properties.

2. Lewis Dot Structures
Here are the Lewis dot structures for some elements. The transition metals, lanthanides and actinides are not displayed by Lewis Dot structures because they do not follow the octet rule. Their first outer orbital shell only has the capacity for two electrons.

3. Lewis Dot Structure Example:
Sulfur has 16 electrons and on the Regents Tables you can see the electron configuration is There are 6 electrons in the outer shell of sulfur.

4. Lewis Dot Structure
This is what sulfur looks like according to the Lewis Dot Diagram: S

5. Lewis Dot Structure
Now it’s your turn to try and draw some elements using the Lewis Dot Structure Potassium Germanium Phosphorus 4. Neon 5. Aluminum

6. Lewis Dot Structure

7. Four Types of Chemical Bonds
Ionic
Polar covalent
Nonpolar covalent
Metallic

8. Ionic Compounds
Ionic compounds are composed of both metals and nonmetals. The bond that is formed is based on electrostatic forces between negatively(anion) and positively(cation) charged ions.
Electrostatic forces remember are the attractions between positive and negative ions that hold the bond together.
Ionic bonding occurs by the transfer of one or more electrons from one atom to another

9. Properties usually solid at room temperature have high melting points
usually do not conduct electricity as a solid
usually dissolve in water
usually conduct electricity when in solution or molten state
1+2-Strong bonds are formed so it is hard to break them apart
3-ions can’t move, and when compound is formed it doesn’t want to give away any electrons
4-since water is a polar molecule it attracts ions

10. Ionic Bonds
An ion is an atom or group of atoms that have a charge. Atoms normally have a neutral charge because most often they have the same number of electrons and protons. They become ions by the loss or addition of one or more electrons. This process is called ionization.
To understand ionic bonding we will develop an understanding of ions.
An ion that has more electrons than protons is called an anion, and an ion that has fewer electrons than protons is called a cation.

11. Ionic Bonds
The interaction of ionic bonds is when atoms gain or lose electrons until the outer shell of electrons is full and stable with 8 electrons. This is part of the octet rule.
Recall octet rule:
When atoms combine to form molecules they generally each lose, gain, or share valence electrons until they attain or share eight and reach a noble gas electron configuration
Stress that the octet rule is useful but it does not work in every case, and there are exceptions to the rule.

12. Ionic Bonds
The number of electrons the atom gains, loses, or shares is called its Valence.
Nonmetals usually have four or more electrons in their outer shell. To make their outer shell full, it’s easier(it takes less ionization energy) for them to gain three or four electrons than to lose four or five electrons.
When you look at the metals, they usually have three or less electrons in their outer shell. Opposite from nonmetals, it is easier for metals to lose three or less electrons than to gain four or more.
Therefore it makes sense that metals and nonmetals bond together easily.

13. Formation of NaF

14. Formation of MgO

15. Ionic Bonds in a crystal of NaCl(S)

16. Covalent Compounds
Covalent compounds are made up of two nonmetals. A single covalent bond is formed when a pair of electrons is shared between two atoms
There are two types of covalent bonds: non-polar covalent and polar covalent.

17. Formation of H2 Nonpolar covalent equal sharing of electrons

18. Formation of I2

19. Formation of HCl polar covalent bonding unequal sharing of electrons

20. Polar Bond vs. NonPolar Bond
The nonpolar covalent bond has an equal sharing of electrons, while the nonpolar covalent bond has an unequal distribution of charge.
A ΔEN of 0 indicates a nonpolar covalent bond.

21. H2O has 2 polar covalent bonds
Polar molecules must have polar bonds, and they must be spaced so the molecule has an uneven distribution of charge.

22. Example-HF H F F H
e- rich
e- poor d+ d-

23. Properties
simple molecular substances have low melting and boiling points
larger more complex compounds will have higher melting and boiling points
usually do not conduct electricity as a solid or when molten or in solution
usually do not dissolve in water
1-weak intermolecular bonds
2-many strong covalent bonds
3-no mobile electrons, no way to transfer electrons
4-molecules are not charged, thus polar water
molecules do not attract them

24. Non-Polar Covalent Bonding
Accounts for the bond that keeps two atoms of the same element together. (Cl2, H2) diatomic elements!
Atoms share electrons from ½ filled orbitals to achieve noble gas configurations
Shared electrons are attracted to both nuclei, which keeps atoms together
Electrons involved in bonding are called shared electron pairs, ones that are not are called lone electron pairs

25. Polar Covalent Bonds Account for the bonding found in HF
One electron from each. atom is shared but not equally due to unequal attraction for shared electrons
The bond is referred to as polar because 2 poles are formed
(+ and -)
Electronegativity values allows us to determine which atom has a greater pull
The atom with the greater electronegativity becomes the negative end of the polar bond.
The atom with the lower electronegativity becomes the positive end of the polar bond

26. Electronegativity
Electronegativity is the tendency of an atom to draw or attract the electrons in a bond toward itself
Electronegativity is like a game of tug-of-war, atom's ability to pull determines what kind of bond it forms
To form a covalent bond, two or more atoms with similar electronegativities will share electrons
Values fall between a low of 0.7 for Fr and a high 4.0 for F
The greater the difference in electronegativity the more polar the bond.

27. > between 1.7 and

29. Double Covalent Bonds
Compounds sometimes share two pairs of electrons and form a double bond.
Examples: O2, CO2

30. Triple Covalent Bond Same idea as single and double
Two atoms of the same element or two different elements share three pairs of electrons and form a triple bond
Example: N2

31. Methods to Classify Bond Type
Non polar covalent bonds=no difference in electronegativity
Polar covalent bonds=difference less than 1.7
Ionic Bonds= difference of 1.7 or more. _ +

32. Using Electronegativity
Use the electronegativity difference to classify these compounds as covalent or ionic:
NH3
NaCl O2
BaCl2
H2O
BaO
Polar COVALENT
IONIC
Nonpolar COVALENT
IONIC
Polar COVALENT
IONIC

33. Metallic Bonding
Valence electrons in metals are free to move throughout the entire substance. Metallic bonds are positive nuclei surrounded by a sea of mobile electrons.
The valence electrons of metals are said to be delocalized.
Delocalized electrons cause metals to conduct electricity very well. Delocalized electrons also explain metallic luster, malleability, and ductility.

34. SEA OF MOBILE VALENCE ELECTRONS
Metallic Bonding
SEA OF MOBILE VALENCE ELECTRONS

35. Energy and Chemical Bonds
Potential energy is stored in chemical bonds.
Energy is released when bonds are formed.
Energy is absorbed to break a chemical bond.

36. Part 2 Naming compounds

37. Common Names
A lot of chemicals have common names as well as the proper IUPAC name.
Chemicals that should always be named by common name and never named by the IUPAC method are:
H2O water, not dihydrogen monoxide
NH3 ammonia, not nitrogen trihydride

38. Formulas of Ionic Compounds
Formulas of ionic compounds are determined from the charges on the ions
atoms ions
    –
Na  +  F :  Na : F :  NaF
   
sodium + fluorine sodium fluoride
Charge balance: = 0

39. Writing a Formula
Write the formula for the ionic compound that will form between Ba2+ and Cl.
Solution:
1. Balance charge with + and – ions
2. Write the positive ion of metal first, and the
negative ion Ba Cl Cl
3. Write the number of ions needed as
subscripts BaCl2

40. Learning Check
Write the correct formula for the compounds containing the following ions:
1. Na+, S2-
a) NaS b) Na2S c) NaS2
2. Al3+, Cl-
a) AlCl3 b) AlCl c) Al3Cl
3. Mg2+, N3-
a) MgN b) Mg2N3 c) Mg3N2

41. CaCl2 = calcium chloride
Naming Compounds
Binary Ionic Compounds:
1. Cation first, then anion
2. Monatomic cation = name of the element
Ca2+ = calcium ion
3. Monatomic anion =root + -ide Cl- = chloride
CaCl2 = calcium chloride

42. Naming Binary Ionic Compounds
Examples:
NaCl
ZnI2
Al2O3
sodium chloride
zinc iodide
aluminum oxide

43. Transition Metals
Elements that can have more than one possible charge MUST have a Roman Numeral to indicate the charge on the individual ion.
1+ or or 3+
Cu+, Cu Fe2+, Fe3+
copper(I) ion iron(II) ion
copper (II) ion iron(III) ion

44. FeCl3 (Fe3+) iron (III) chloride
Names of Variable Ions
These elements REQUIRE Roman Numerals because they can have more than one possible charge:
anything except Group 1A, 2A, Ag, Zn, Cd, and Al
(You should already know the charges on these!)
Or another way to say it is: Transition metals and the metals in groups 4A and 5A (except Ag, Zn, Cd, and Al) require a Roman Numeral.
FeCl3 (Fe3+) iron (III) chloride
CuCl (Cu+ ) copper (I) chloride
SnF (Sn4+) tin (IV) fluoride
PbCl (Pb2+) lead (II) chloride

45. Learning Check
Complete the names of the following binary compounds with variable metal ions:
FeBr2 iron (_____) bromide
CuCl copper (_____) chloride
SnO2 ___(_____ )

46. Polyatomic Ions
NO3-
nitrate ion
NO2-
nitrite ion

47. Polyatomic Ions
Use Table E

48. Ternary Ionic Nomenclature
Writing Formulas
Write each ion, cation first. Don’t show charges in the final formula.
Overall charge must equal zero.
If charges cancel, just write symbols.
If not, use subscripts to balance charges.
Use parentheses to show more than one of a particular polyatomic ion.
Use Roman numerals indicate the ion’s charge when needed (stock system)

49. Ternary Ionic Nomenclature
Sodium Sulfate
Na+ and SO4 -2
Na2SO4
Iron (III) hydroxide
Fe+3 and OH-
Fe(OH)3
Ammonium carbonate
NH4+ and CO3 –2
(NH4)2CO3

50. Learning Check 1. aluminum nitrate a) AlNO3 b) Al(NO)3 c) Al(NO3)3
2. copper(II) nitrate
a) CuNO3 b) Cu(NO3)2 c) Cu2(NO3)
3. Iron (III) hydroxide
a) FeOH b) Fe3OH c) Fe(OH)3
4. Tin(IV) hydroxide
a) Sn(OH)4 b) Sn(OH)2 c) Sn4(OH)

51. Naming Ternary Compounds
Contains at least 3 elements
There MUST be at least one polyatomic ion
(it helps to circle the ions)
Examples:
NaNO3 Sodium nitrate
K2SO4 Potassium sulfate
Al(HCO3)3 Aluminum bicarbonate

52. Learning Check Match each set with the correct name:
Na2CO3 a) magnesium sulfite
MgSO3 b) magnesium sulfate
MgSO4 c) sodium carbonate
2 . Ca(HCO3)2 a) calcium carbonate
CaCO3 b) calcium phosphate
Ca3(PO4)2 c) calciumbicarbonate

53. Mixed Practice!
Name the following:
Na2O
CaCO3
PbS2
Sn3N2
Cu3PO4
HgF2

54. Mixed Up… The Other Way Write the formula: Copper (II) chlorate
Calcium nitride
Aluminum carbonate
Potassium bromide
Barium fluoride
Cesium hydroxide

55. Naming Molecular Compounds
All are formed from two or more nonmetals.
Naming Molecular Compounds
CO2 Carbon dioxide
Ionic compounds generally involve a metal and nonmetal
(NaCl)
BCl3 boron trichloride
CH4 methane

56. Molecular (Covalent) Nomenclature for two nonmetals
Prefix System (binary compounds)
1. Less electronegative atom comes first.
2. Add prefixes to indicate # of atoms. Omit mono- prefix on the FIRST element. Mono- is OPTIONAL on the SECOND element (in this class, it’s NOT optional!).
3. Change the ending of the second element to -ide.

57. Molecular Nomenclature Prefixes
mono-
di-
tri-
tetra-
penta-
hexa-
hepta-
octa-
nona-
deca-
NUMBER 1 2 3 4 5 6 7 8 9 10

58. Molecular Nomenclature: Examples
CCl4
N2O
SF6
carbon tetrachloride
dinitrogen monoxide
sulfur hexafluoride

59. More Molecular Examples
arsenic trichloride
dinitrogen pentoxide
tetraphosphorus decoxide
AsCl3
N2O5
P4O10

60. Learning Check
Fill in the blanks to complete the following names of covalent compounds.
CO carbon ______oxide
CO2 carbon _______________
PCl3 phosphorus ______chloride
CCl4 carbon ________chloride
N2O _____nitrogen _____oxide

61. Learning Check P2O5 a) phosphorus oxide Cl2O7 a) dichlorine heptoxide
b) phosphorus pentoxide
c) diphosphorus pentoxide
Cl2O7
a) dichlorine heptoxide
b) dichlorine oxide
c) chlorine heptoxide
Cl2 a) chlorine
b) dichlorine
c) dichloride

62. Name the following compounds:
1. CaO
a) calcium oxide
b) calcium(I) oxide c) calcium (II) oxide
2. SnCl4
a) tin tetrachloride
b) tin(II) chloride
c) tin(IV) chloride
3. N2O3
a) nitrogen oxide
b) dinitrogen trioxide
c) nitrogen trioxide

63. Solution Name the following compounds: 1. CaO 2. SnCl4
3. N2O3
a) calcium oxide
c) tin(IV) chloride
b) Dinitrogen trioxide

64. Mixed Practice Dinitrogen monoxide Potassium sulfide
Copper (II) nitrate
Dichlorine heptoxide
Chromium (III) sulfate
Iron (III) sulfite
Calcium oxide
Barium carbonate
Iodine monochloride

65. Mixed Practice BaI2 P4S3 Ca(OH)2 FeCO3 Na2Cr2O7 I2O5 Cu(ClO4)2 CS2
B2Cl4

66. Acid Nomenclature USE TABLE K
Acids
Compounds that form H+ in water.
Formulas usually begin with ‘H’.
In order to be an acid instead of a gas, binary acids must be aqueous (dissolved in water)
Ternary acids are ALL aqueous
Examples:
HCl (aq) – hydrochloric acid
HNO3 – nitric acid
H2SO4 – sulfuric acid

67. Name the following acids
HI (aq)
HCl
H2SO3
HNO3
HIO4

68. Write the Formula! Hydrobromic acid Nitrous acid Carbonic acid
Phosphoric acid

69. Part 3 Forces Between Molecules

70. Opposites attract H-Cl H is the positive end of the molecule
Cl is more negative (has more electrons)

71. Same charges repel
H2 molecules repel eachother, it explains why they are gases.
Hydrogen is a nonpolar molecule…both ends are the same, positive.

72. Same charges still repel
All the nitrogens are surrounded by a lot of electrons…each is slightly negative.

73. Hydrogen Bonding Hydrogen bonding explains many of waters properties.
High boiling point
Explains why ice floats
Explains why salts dissolve

74. Hydrogen bonds are strong forces of attraction.

75. Ionic Solids Dissociate in Water

77. + - O H + - Property Ionic Covalent (simple) Metallic Bonding Type
Which types of atoms does it involve?
How are they structured?
An example of the bonding type is
What is the melting point or boiling point like?
Are they magnetic?
Are they soluble in water?
Do they conduct electricity when solid liquid, gas or dissolved?
Give an example substance with a use
Bonding Type
Metal and Non-metal atoms
Non-metal atoms only
Metal atoms only
Form giant lattices that are brittle so are easy to crush
Form molecules with weak forces of attraction so are mainly gases and liquids
Forms a strong lattice structure + - O H Na + Cl -
High melting point and boiling point
Low melting point and boiling point
High melting point and boiling point
Not magnetic
Not magnetic
Some are magnetic
Many are soluble in water
Are not soluble in water
Are not soluble in water
Will conduct electricity when molten or dissolved in water
Will conduct electricity when solid o molten
Will not conduct electricity
Copper
Used for wiring electrical appliances
Sodium chloride
Used to flavour food
Water
Needed for life

78. Bonding Type Property Ionic Covalent (simple) Metallic
Which types of atoms does it involve?
How are they structured?
An example of the bonding type is
What is the melting point or boiling point like?
Are they magnetic?
Are they soluble in water?
Do they conduct electricity when solid liquid, gas or dissolved?
Give an example substance with a use