OCTET RULE & IONS

• Most elements, except noble gases, combine to form compounds. Compounds are the result of the formation of chemical bonds between two or more different elements.
• In the formation of a chemical bond, atoms lose, gain or share valence electrons to complete their outer shell and attain a noble gas configuration.
• This tendency of atoms to have eight electrons in their outer shell is known as the octet rule.

Formation of Ions:

• An ion (charged particle) can be produced when an atom gains or loses one or more electrons.
  

  

• A cation (+ ion) is formed when a neutral atom loses an electron.
• An anion (- ion) is formed when a neutral atom gains an electron.

Cl atom

ion

IONIC CHARGES

• The ionic charge of an ion is dependent on the number of electrons lost or gained to attain a noble gas configuration.
• For most main group elements, the ionic charges can be determined from their group number, as shown below:

		Metals Lose Valence Electrons			Nonmetals Gain Valence Electrons
Noble Gases		1A (1)	2A (2)	3A (13)	5A (15)	6A (16)	7A (17)		Noble Gases
He
Ne									Ne
Ar						S			Ar
Kr									Kr
Xe									Xe

• All other ionic charges need to be memorized and known in order to write correct formulas for the compounds containing them.

COMPOUNDS

• Compounds are pure substances that contain 2 or more elements combined in a definite proportion by mass.
• Compounds can be classified as one of two types: Ionic and molecular (covalent)
  

  

• Ionic compounds are formed by combination of a metal and a non-metal. The smallest particles of ionic compounds are ions.
• Molecular compounds are formed by combination of 2 or more non-metals. The smallest particles of molecular compounds are molecules.
• The nature and type of the chemical bond is directly responsible for many physical and chemical properties of a substance: (e.g. melting point, conductivity)
  

  

• When the conductivity apparatus is placed in salt solution, the bulb will light. But when it is placed in sugar solution, the bulb does not light.
• This difference in conductivity between salt and sugar is due to the different types of bonds between their atoms.

IONIC COMPOUNDS

• Ionic compounds contain ionic bonds, which occur when electrons are transferred between two atoms.
  

  

• Ionic bonds occur between metals and non-metals.
• Atoms that lose electrons (metals) form positive ions (cations).
• Atoms that gain electrons (non-metals) form negative ions (anions).
• The smallest particles of ionic compounds are ions (not atoms).

(a)

(b)

(c)

IONIC CHARGES & FORMULAS

• The formula of an ionic compound indicates the number and kinds of ions that make up the ionic compound.
• The sum of the ionic charges in the formula is always zero, which indicates that the total number of positive charges is equal to the total number of negative charges.
• For example, the +1 charge on the sodium ion is cancelled by the -1 charge on the chloride ion, to form a net zero charge.
  

  

• When charges between the two ions do not balance, subscripts are used to balance the charges.
• For example, since each magnesium loses 2 electrons, and each chloride gains one electron, 2 chlorides are needed to balance the charge of the magnesium ion. Therefore, magnesium chloride is written as .
  

  

NAMING & WRITING IONIC FORMULAS

• When writing ionic formula, knowing the charge of the ions are important since the net charge on the compound must be zero.
• Some elements produce only one ion (Type I) while others produce two or more ions (Type II).
• Differentiating between type I and II ions is important, since the naming system is different for each. Shown below are the common ions of each type:

• Note that most main-group elements are type I (except Sn and Pb), and most transition elements are type II (except Ag, Cd and Zn).

NAMING & WRITING IONIC FORMULAS

Binary Ionic Compounds (Type I):

• Binary compounds contain only two elements.
• Type I are those cations that form only one ion.
• In these compounds, charges of the cations must equal the charges of the anions since the net charge is zero.
• Subscripts are used to balance the charges between cations and anions.

• Name the cation first, the anion last.
• The cation name is the same as the name of the metal it forms from.
• The anion name takes the root of non-metal and the ending "-ide".

Examples:

1. Write formulas for the following ionic compounds:
   calcium chloride:
   sodium sulfide:
2. Name the following ionic compounds:
   :
   :

NAMING & WRITING IONIC FORMULAS

Binary Ionic Compounds (Type II):

• Type II ions are those cations that form more than one ion.
• When naming compounds formed from these ions, include the ionic charge as Roman numeral, in parentheses, after the metal's name.
• This method of nomenclature is called the "stock" system.

Examples:

1. Name each of the following compounds:
   :
   :
   
2. Write formulas for each of the following compounds:
   tin(II) bromide:
   titanium (IV) oxide:
   chromium (III) oxide

COVALENT COMPOUNDS

• Covalent bonds occur when electrons are shared between two atoms.
  

  

• Covalent bonds occur between two non-metals.
• The smallest particle of a covalent compound is a molecule.
• Two types of covalent bonds exist: polar and nonpolar.
• Nonpolar covalent bonds occur between similar atoms. In these bonds the electron pair is shared equally between the two protons.
  

  

• Polar covalent bonds occur between different atoms. In these bonds the electron pair is shared unequally between the two protons. As a result there is a charge separation in the molecule, and partial charges on each atom.
  

  

Examples:

NAMING & WRITING COVALENT FORMULAS

Binary Covalent Compounds:

• These compounds are named similar to ionic compounds, with the second element named based on its root and suffix "-ide".
• Greek prefixes are used to indicate the number of atoms in these compounds:

• The first atom uses a prefix only when more than one atom is present.
• The second atom always uses a prefix.

Examples:

1. Name the following molecular compounds:

SUMMARY OF NAMING BINARY COMPOUNDS

POLYATOMIC IONS

• Some ionic compounds contain polyatomic ions, an ion composed of several atoms bound together.
• Some common polyatomic ions are:

• When writing formulas for compounds containing polyatomic ions, treat the polyatomic ion as one group.

Examples:

1. Write formulas for the following polyatomic compounds:
   sodium carbonate:
   ammonium sulfide:
   magnesium bicarbonate:

POLYATOMIC IONS

• Polyatomic ionic compounds are named by naming the cation first, followed by the polyatomic ion.

Examples:

2. Name the following polyatomic compounds:
   :

NAMING ACIDS

• Acids are molecular compounds that form ions when dissolved in water.

Naming Binary Acids:

• Formulas are written similar to binary ionic compounds, assigning a +1 charge to hydrogen.
• When naming the acids, use hydro- prefix, followed by the name of the non-metal with an -ic ending, followed with the word acid.

Naming Polyatomic Acids:

• Several polyatomic acids are important in the study of chemistry, and their names must be learned. These acids and the polyatomic ions that form from their ionization are listed below:

ELECTRONEGATIVITY

• Electronegativity (E.N.) is the ability of an atom involved in a covalent bond to attract the bonding electrons to itself.
• Linus Pauling derived a relative Electronegativity Scale based on Bond Energies:
  

  

Electronegativities of the Elements

BOND POLARITY & ELECTRONEGATIVITY

Examples:

1. Hydrogen molecule

2. Hydrogen chloride

SUMMARY OF BONDING

LEWIS STRUCTURES

• Lewis structures use Lewis symbols to show valence electrons in molecules and ions of compounds.

• In a Lewis structure, a shared electron pair is indicated by two dots between the atoms, or by a dash connecting them.
  
• Unshared pairs of valence electrons (called lone pairs) are shown as belonging to individual atoms or ions.
• Writing correct Lewis structures for covalent compounds requires an
  
  
  understanding of the number of bonds normally formed by common nonmetals.

4A	5A	6A	7A	8A
- X.	.	.	: .
4 bonds	3 bonds	2 bonds	1 bond	0 bonds
			:	:Noe:

LEWIS STRUCTURES

• Covalent structures are best represented with electron-dot symbols or Lewis structures. Structures must satisfy octet rule (except hydrogen).
  

  

covalent bond

• More complex Lewis structures can be drawn according to the following steps:

1. Count the total number of valence electrons to be used in the structure. For ions, add one electron for each negative charge, and subtract one electron for each positive charge.
2. Write a skeleton structure, arranging the atoms in the most symmetrical pattern. Remember the number of bonds commonly formed by atoms.
3. Connect each atom by a dashed line representing 2 electrons each.
4. Determine the number of electrons left by subtracting number of electrons used from the total determined in step 1 .
5. Distribute the remaining electrons around the atoms in pairs in order to achieve eight electrons around each (octet rule). Hydrogen is an exception (doublet).
6. If there are not enough electrons to achieve octet rule for each atom, form double or triple bonds by sharing more than one pair of electrons.

LEWIS STRUCTURES

Examples:

1. water
2. carbon dioxide ( )
3. carbonate ion ( )
4. ammonia
5. chlorate ion

EVALUATING LEWIS STRUCTURES

• When evaluating Lewis structures for correctness, two points must be considered:

1. Are the correct number of electrons present in the structure?
2. Is octet rule satisfied for all elements? (Hydrogen is an exception)

Examples:

1. 
   
2. 
   

   

MOLECULAR SHAPES

• The three-dimensional shape of the molecules is an important feature in understanding their properties and interactions.
• All binary molecules have a linear shape since they only contain two atoms.
• More complex molecules can have various shapes (linear, bent, etc.) and need to be predicted based on their Lewis structures.
• A very simple model, VSEPR (Valence Shell Electron Pair Repulsion) Theory, has been developed by chemists to predict the shape of large molecules based on their Lewis structures.
• Based on VSEPR, the electron pair groups in a molecule will repel one another and seek to minimize their repulsion by arranging themselves around the central atom as far apart as possible.
• Electron pair groups can be defined as any one of the following: bonding pairs, non-bonding pairs and multiple bonds.
• A summary of VSEPR predictions are listed below:

MOLECULAR SHAPES

• Examples of various shapes are shown below.

Linear :

Bond angle:

Trigonal Planar ( ):

Bond angle:
Polarity:

Bent (AX2E):

Bond angle:

Tetrahedral ( ):

Bond angle:
Polarity:
Exceptions:

MOLECULAR SHAPES

Pyramidal ( ):
Bond angle:
Polarity:

Bent ( ):
Bond angle:

Summary of Shapes:

• Symmetrical shapes that have polar bonds, but are usually non-polar:

• Unsymmetrical shapes that have polar bonds, but are always polar:

ATTRACTOVE FORCES IN MOLECULES

• Many differences in the properties of solids and liquids are associated with the attractive forces that keep their molecules together. These forces range from very strong to very weak and are responsible for the wide range of properties observed in these substances.
• Ionic solids have high melting points that are due to the strong nature of the ionic bond and the forces of attraction between the cation and the anion.
  

  

  
  

  

• Covalent molecules also contain attractive forces, but of a weaker nature than the ionic compounds. The forces between covalent molecules can be divided into three types:

1. Dispersion Forces:

These forces exist between non-polar molecules, and are the weakest of the 3 types of covalent forces. These forces are caused by temporary shifts in distribution of electrons in a non-polar molecule that cause a temporary dipole.

2. Dipole-Dipole Attractions:

These forces exist in polar molecules and are moderate strength. The negative end of a polar molecule is attracted to the positive end of another such molecule by these forces.

3. Hydrogen Bonds:

These are the strongest of the 3 types of forces and are formed between molecules that have a hydrogen attached to the very electronegative elements (F, O and N).

ATTRACTOVE FORCES IN MOLECULES

• Hydrogen bonding in water is responsible for its many unique properties and behavior in the body. Two representations of this bonding is shown below:
  

  

  
  

  

ATTRACTOVE FORCES IN MOLECULES

Examples:

1. Indicate the major type of molecular forces expected of each of the following:
   a)
   b)
   c)
2. Identify the strongest attractive forces between molecules of each of the following:
   a)
   b)
   c) HCl
   d)
3. Identify the substance in each pair that would have the higher boiling point:
   a) HF or HBr
   b) or
   c) or