1. Comparison of Properties Ionic Compounds Covalent Compounds Metals
Chemical Bonding
Comparison of Properties
Ionic Compounds
Covalent Compounds
Metals

2. Essential Questions
Why/How do atoms combine with one another to form the vast array of chemical substances that exist?
What is ionic, covalent and metallic bonding and how do the types of bonding determine properties of matter? 

3. Properties of Matter
Macroscopic properties of matter vary greatly due to the type of bonding

4. What is a chemical bond?
An attractive force that holds two atoms together
Can form by
The attraction of positive ion to a negative ion or
The attraction of the positive nucleus of one atom and the negative electrons of another atom

5. Bond
the interaction between two or more atoms that allows them to form a substance different from the independent atoms.
involves the outer (valence) electrons of the atoms.
These electrons are
transferred from one atom to another or shared between them.

6. Chemical Bond Energy Considerations
A chemical bond forms when it is energetically favorable
when the energy of the bonded atoms is less than the energies of the separated atoms.
Al + I2

7. Bonding
Chemical compounds are formed by the joining of two or more atoms.
A stable compound occurs when the total energy of the combination has lower energy than the separated atoms.
The bound state implies a net attractive force between the atoms ... a chemical bond.

8. Energy Changes in Bonding
When bonds are formed, energy is released.
Demonstrations:
Formation of an Ionic Compound: Mg + O2
Formation of a Molecular Compound: S + O2

9. Breaking Bonds
In order to break bonds energy must be added, usually in the form of heat, light, or electricity.
Demonstration: Electrolysis of water
Demo: Decomposition of Nitrogen Triiodide

10. Three Types of Bonding
Ionic
Metallic
Covalent

11. Chemical Bonds In chemical bonds, atoms can either transfer or
share their valence electrons.

12. When atoms transfer electrons Ionic Bonds
When one or more atoms lose electrons and other atoms gain them in order to produce a noble gas electron configuration, the bond is called an ionic bond.

13. Ionic Bonding metallic atoms tend to lose electrons
When they do so, they become positively charged ions which are called cations.
Nonmetallic atoms tend to gain electrons to become negatively charged ions which are called anions.
These oppositely charged cations and anions are attracted to one another because of their opposite charges.
That attraction is called an  ionic bond. We often refer to the charge on the ion as the oxidation state of that element.

14. Positive Ion (Cation) Formation Negative Ion (Anion) Formation
Na has one valence electron.
It loses it to Chlorine.
Na now has a filled valence shell. (an octet)
Becomes positive one in charge
Chlorine has seven valence electrons.
It gains one electron from Na.
Chlorine now has filled octet.
Chlorine has a negative one charge. (Chloride ion)
Na+1 attracts Cl-1 and forms the ionic bond.

15. Metals
Nonmetals

16. Ionic Bonds Part 1 Part 2 Part 3
Part 2
Part 3

17. Ion Formation
All of the elements in Group I have one electron in their outermost energy level.
All of these elements can lose that one valence electron.
These atoms become cations with a positive one charge.

18. Elements in Group II have two electrons in their outermost energy level.
So, when these elements lose electrons, they lose two electrons and take on a positive two charge.

19. The transition metals and the metals to the right of them generally form more than one ion.
We call these elements multivalent. The charges on their ions are not always predictable, although some patterns do exist.
A few of the transition elements form only one ion or oxidation state. For example zinc ion, silver ion and scandium ion.
Zn2+ zinc ion
Ag+ silver ion
Sc3+ scandium ion

20. Anions Nonmetals tend to gain electrons.
The halogens - fluorine, chlorine, bromine, and iodine - have a strong attraction for electrons.
Their outermost energy levels are almost full. There is only room for one more electron in the outer energy levels for each of those atoms. Consequently, the elements fluorine, chlorine, bromine, and iodine will gain one electron, and become anions with a negative one charge.
Oxygen, sulfur, and the other elements in that family will gain two electrons.
In the next group over, nitrogen, phosphorus and arsenic can take on three electrons.

21. +1 +2 +3 -3 -2 -1

22. Ionic Nomenclature Naming Ionic Compounds Video of the Process

23. Ionic Compounds Made of cations and anions Metals and nonmetals
The electrons lost by the cation are gained by the anion
The cation and anions surround each other
Smallest ratio of ions in an ionic compound is a FORMULA UNIT.

24. K+1 Ca+2 Cations Positive ions Formed by losing electrons
More protons than electrons
usually Metals
K+1
Has lost one electron
Ca+2
Has lost two electrons

25. F-1 O-2 Anion A negative ion Has gained electrons Non metals
Charge is written as a super script on the right.
F-1
Has gained one electron
O-2
Has gained two electrons

26. Formula Unit
The smallest whole number ratio of atoms in an ionic compound.
Ions surround each other so you can’t say which is hooked to which

27. Naming Ions We will use the systematic way
Cation- if the charge is always the same just write the name of the metal
Transition metals can have more than one type of charge
Indicate the charge with a Roman numeral in parentheses

28. Name these
Na+1
Ca+2
Al+3
Fe+3
Fe+2
Pb+2
Li+1

29. Write Formulas for these
Potassium ion
Magnesium ion
Copper (II) ion
Chromium (VI) ion
Barium ion
Mercury (II) ion

30. Naming Anions
Change the element ending to – ide
F-1 Fluorine

31. Name these
Cl-1
N-3
Br-1
O-2
Ga+3

32. Write these
Sulfide ion
iodide ion
phosphide ion
Strontium ion

33. Polyatomic ions Groups of atoms that stay together and have a charge
You must memorize these or use an ion sheet… common examples
Acetate C2H3O2-1
Nitrate NO3-1
Nitrite NO2-1
Hydroxide OH-1
Permanganate MnO4-1
Cyanide CN-1

34. More Polyatomic ions Sulfate SO4-2 Sulfite SO3-2 Carbonate CO3-2
Chromate CrO4-2
Dichromate Cr2O7-2
Phosphate PO4-3
Phosphite PO3-3
Ammonium NH4+1

35. Practice with Ions
Use the practice worksheet to determine the ions formed.
Learn to use your periodic table and pink sheet to determine charges (oxidation state.)

36. Binary Ionic Compounds
Binary Compounds
2 elements.
a cation and an anion.
To write the names just name the two ions.
Easy with Representative elements
Groups 1, 2, 13
NaCl = Na+ Cl- = sodium chloride
MgBr2 = Mg+2 Br- = magnesium bromide

37. Naming Binary Ionic Compounds with Variably Charged Cations
The problem comes with the transition metals (Groups 3-12) since their charge can vary
Need to figure out their charges
The compound must be neutral
same number of + and – charges.
Use the anion to determine the charge on the positive ion
Charge of the cation is a Roman numeral in the name

38. Example Write the name of CuO Need the charge of Cu O is -2
copper must be +2
Copper (II) chloride

39. Example Name CoCl3 Cl is -1 and there are three of them = -3
Co must be +3 Cobalt (III) chloride

40. Another Example Write the name of Cu2S.
Since S is -2, the Cu2 must be +2, so each one is +1.
copper (I) sulfide

41. Last Example Fe2O3 Each O is -2 3 x -2 = -6
3 Fe must = +6, so each is +2.
iron (III) oxide

42. Naming Binary Ionic Compounds
Write the names of the following
KCl
Na3N
CrN
Sc3P2
PbO
PbO2
Na2Se

43. Ternary Ionic Compounds
Will have polyatomic ions
At least three elements
Name the ions
NaNO3
CaSO4
CuSO3
(NH4)2O

44. Ternary Ionic Compounds
LiCN
Fe(OH)3
(NH4)2CO3
NiPO4

45. Writing Formulas Given the name write the formula
The charges have to add up to zero
Write down each ion with charges
Make the charges equal by adding subscripts
Put polyatomic ions in parentheses if you need more than one of them

46. Writing Formulas Example
Write the formula for calcium chloride.

47. Another Example
Aluminum nitrate

48. Write the formulas for these
Lithium sulfide
tin (II) oxide
tin (IV) oxide
Magnesium fluoride
Copper (II) sulfate
Iron (III) phosphide

49. Write the formulas for these
gallium nitrate
Iron (III) sulfide
Ammonium chloride
ammonium sulfide
barium nitrate

50. Things to look for
If cation has (Roman Numeral), the number is the charge
If anions end in -ide they are probably off the periodic table (Monoatomic)
If anion ends in -ate or -ite it is polyatomic

51. Ionic Solids
Ionic solids are solids composed of ionic particles (ions).
These ions are held together in a regular array by ionic bonding.
Ionic bonding results from attractive interactions from oppositely charged ions.
In a typical ionic solid, positively charged ions are surrounded by negatively charged ions and vice-versa.
The close distance between these oppositely charged particles results in very strong attractive forces.
The alternating pattern of positive and negative ions continues in three dimensions.
The regular repeating pattern is analogous to the tiles on a floor or bricks on a wall.
called the crystal lattice.

52. Ionic Compounds Crystalline solids (made of ions)
High melting and boiling points
Conduct electricity when melted or dissolved in water
Demo: Electrolytes
Many are soluble in water but not in non-polar liquid

53. Comparison of Conductivity

54. Common Ionic Compounds
NaCl - sodium chloride - table salt
KCl - potassium chloride - present in "light" salt (mixed with NaCl)
CaCl2 - calcium chloride - driveway salt
NaOH - sodium hydroxide - found in some surface cleaners as well as oven and drain cleaners
CaCO3 - calcium carbonate - found in calcium supplements
NH4NO3 - ammonium nitrate - found in some fertilizers

55. Ionic vs Molecular

56. Covalent (Molecular) Compounds
Gases, liquids, or solids (made of molecules)
Low melting and boiling points
Poor electrical conductors in all phases
Many soluble in non-polar liquids but not in water

57. Molecular (Covalent) Substances

58. Covalent Network Solids
Covalent because combinations of nonmetals
Interconnected
very hard and brittle
Insoluble
Extreme melting and boiling points
Diamond

59. Covalent Bonds
involve the sharing of a pair of valence electrons by two atoms
Such bonds lead to stable molecules if they share electrons in such a way as to create a noble gas configuration for each atom

60. Covalent bonding can be visualized with the aid of a Lewis Structure

61. Polar Covalent Bonds
Covalent Bonds in which the sharing of the electron pair is unequal
the electrons spend more time around the more nonmetallic atom
In such a bond there is a charge separation with one atom being slightly more positive and the other more negative……. will produce a dipole moment.

62. Types of Covalent bonds
Pure Covalent (also called non-polar covalent) bonds are ones in which both atoms share the electrons evenly
By evenly, we mean that the electrons have an equal probability of being at a certain radius from the nuclei of either atom.
Polar covalent bonds are ones in which the electrons have a higher probability of being in the proximity of one of the atoms
Determined by Electronegativity Difference

63. Electronegativity
the periodic property that indicates the strength of the attraction an atom has for the electrons it shares in a bond.
Atoms with high electronegativities tend to hold tightly to their electrons or to form negative ions.
These elements are found to the upper right on the periodic table.
Atoms with low electronegativities tend to have a lower attraction for their electrons and may form positive ions.
These elements are found to the lower left on the periodic table.

64. Pure covalent or Non-polar covalent bond
Electronegativity difference of 0.3 or less in between the two atoms.
A pure covalent bond can form between two atoms of the same element (such as in diatomic oxygen molecule)
or atoms of different elements that have similar electronegativies (such as in the carbon and hydrogen atom in methane).

65. Polar Covalent Bond
A is a pair of electrons shared between two atoms with significantly different electronegativities (from 0.3 to 1.7 difference).
These bonds tend to form between highly electronegative non-metals and other non-metals, such as the bond between hydrogen and oxygen in water.

66. Ionic Bonds
In compounds that have elements with very different electronegativities (greater than 1.7 difference), the electrons can be considered to have been transferred to form ions.

67. Many of the properties of a compound, such as solubility and boiling point, depend, in part, on the degree of the polarity of its bonds.

68. Examples to Determine Bond Character
Using electronegativity in the prediction of the polarity of a chemical bond.
sodium bonded to chlorine
Difference between the electronegativities of Na(0.9) and Cl(3.0) are so great that they form an ionic bond.
The hydrogen molecule (2 H atoms bonded to each other)
zero electronegativity difference, form a non-polar covalent bond.

69. Bond Character Nonpolar-Covalent bonds (H2) Polar-Covalent bonds (HCl)
Electrons are equally shared
Electronegativity difference of 0 to 0.3
Polar-Covalent bonds (HCl)
Electrons are unequally shared
Electronegativity difference between .3 and 1.7
Ionic Bonds (NaCl)
Electrons are transferred
Electronegativity difference of more than 1.7

70. Diatomic Molecules hydrogen gas H2 the halogens:
chlorine Cl2
fluorine F2
bromine Br2
iodine I2
Nitrogen N2
Oxygen O2
Pneumonic Device to remember the diatomic molecules: Professor BrINClHOF

71. Metals and Metallic Bonding
Typical Properties of Metals
Malleable
Ductile
Good Conductors of Heat and Electricity
Generally high melting and boiling points

72. Metallic Bonds
The properties of metals suggest that their atoms possess strong bonds
yet the ease of conduction of heat and electricity suggest that electrons can move freely in all directions in a metal
The general observations give rise to a picture of "positive ions in a sea of electrons" to describe metallic bonding.

73. Metal Properties Malleable and Ductile Strong and Durable
Good conductors of heat and electricity.
Their strength indicates that the atoms are difficult to separate… strong bonds
but malleability and ductility suggest that the atoms are relatively easy to move in various directions.
The electrical conductivity suggests that it is easy to move electrons in any direction in these materials.
The thermal conductivity also involves the motion of electrons. All of these properties suggest the nature of the metallic bonds between atoms. (Electron sea model)

74. Metallic Bonding Electron Sea Model
Explained by the Electron Sea Model
the atoms in a metallic solid contribute their valence electrons to form a “sea” of electrons that surrounds metallic cations.
delocalized electrons are not held by any specific atom and can move easily throughout the solid.
A metallic bond is the attraction between these electrons and the metallic cation.

75. Metallic Bonding the Electron Sea Model
The more delocalized electrons the stronger the bond

76. A mixture of elements that has metallic properties is called an alloy.
Two types of alloys
An interstitial alloy is one in which the small holes in a metallic crystal are filled by other smaller atoms.
A substitutional alloy is one in which atoms of the original metal are replaced by other atoms of similar size.

77. Ionic Compounds
Covalent Compounds
Metallic Compounds
-Formed from a combination of metals and nonmetals.
-Electron transfer from the cation to the anion.
-Opposite charged ions attract each other.
-Formed from a combination of nonmetals.
-Electron sharing between atoms.
-Formed from a combination of metals
-“sea of electrons”;
electrons can move among atoms
Solids at room temperature
Can be solid, liquid, or gas at room temperature.
High melting points
Low melting points
Various melting points
Dissolve well in water
Do not dissolve in water (Sugar is an exception)
Do not dissolve in water.
Conduct electricity only when dissolved in water; electrolytes
Do not conduct electricity; non electrolytes
Conduct electricity in solid form.
Brittle, hard
Soft
Metallic compounds range in hardness. Group 1 and 2 metals are soft; transition metals are hard. Metals are malleable, ductile, and have luster.