IV. Periodicity and Reactivity; Chemical Reactions; Biochemistry and Organic Chemistry 23 QUESTIONS

● Periodicity and chemical reactivity

– Chemical reactivity

– Periodic trends in electron configurations, atomic properties such as radius, electronegativity, ionization potential, and chemical reactivity

– Relationship between bond types and periodicity

● Chemical reactions

– Equation balancing and stoichiometry

– Reaction types

– Reaction mechanisms and kinetics

– Chemical equilibrium

– Redox chemistry and electrochemistry

● Biochemistry and Organic Chemistry

– Organic functional groups and their reactions

– Biologically important compounds and reactions

I. Periodicity and chemical reactivity

A. Chemical Reactivity

Alkali and Alkaline (Groups 1 and 2) earth metal quickly turn into an oxide in nature.  Groups 1 and 2 do not really existed in nature as free metal. Noble gases are nearly chemical inert.  Heaviest noble gases from a few compounds with oxygen and fluorine such as KrF₂ and XeO₄.

Electronegativity

Electronegativity is the power of an atom in a molecule to attract electrons to itself.  Electronegativity increases from left right and from bottom top (decrease from top to bottom). 

·         note hydrogen combining ratios (LIH, BeH₂, BH₃, CH₄, H₃N, H₂O, HF) and acid/base properties of oxides (basic for metals, acidic for non-metals)

Pauling’s scale of electronegativities for 10 common elements with their values:

Valence and Oxidation Numbers

In chemistry, valence electrons are the electrons contained in the outermost, or valence, electron shell of an atom. Valence electrons are important in determining how an element reacts chemically with other elements: The fewer valence electrons an atom holds, the less stable it becomes and the more likely it is to react.

Oxidation Number

1.      Group IA always +1

2.      Group IIA always +2

3.      Group IIIA commonly +3, Al always+3

4.      Group IV +4 or -4 (usually covalent)

5.      Group VA commonly -3

6.      Group VIA commonly -2

7.      Group VIIA commonly -1

By definition, the oxidation number of an atom is the charge that atom would have if the compound was composed of ions.

1. The oxidation number of an atom is zero in a neutral substance that contains atoms of only one element. Thus, the atoms in O₂, O₃, P₄, S₈, and aluminum metal all have an oxidation number of 0.

2. The oxidation number of simple ions is equal to the charge on the ion. The oxidation number of sodium in the Na⁺ ion is +1, for example, and the oxidation number of chlorine in the Cl^- ion is -1.

3. The oxidation number of hydrogen is +1 when it is combined with a nonmetal as in CH₄, NH₃, H₂O, and HCl.

4. The oxidation number of hydrogen is -1 when it is combined with a metal as in. LiH, NaH, CaH₂, and LiAlH₄.

5. The metals in Group IA form compounds (such as Li₃N and Na₂S) in which the metal atom has an oxidation number of +1.

6. The elements in Group IIA form compounds (such as Mg₃N₂ and CaCO₃) in which the metal atom has a +2 oxidation number.

7. Oxygen usually has an oxidation number of -2. Exceptions include molecules and polyatomic ions that contain O-O bonds, such as O₂, O₃, H₂O₂, and the O₂^2- ion.

8. The elements in Group VIIA often form compounds (such as AlF₃, HCl, and ZnBr₂) in which the nonmetal has a -1 oxidation number.

9. The sum of the oxidation numbers in a neutral compound is zero.

H₂O: 2(+1) + (-2) = 0

10. The sum of the oxidation numbers in a polyatomic ion is equal to the charge on the ion. The oxidation number of the sulfur atom in the SO₄^2- ion must be +6, for example, because the sum of the oxidation numbers of the atoms in this ion must equal -2.

SO₄^2-: (+6) + 4(-2) = -2

11. Elements toward the bottom left corner of the periodic table are more likely to have positive oxidation numbers than those toward the upper right corner of the table. Sulfur has a positive oxidation number in SO₂, for example, because it is below oxygen in the periodic table.

SO₂: (+4) + 2(-2) = 0

Metal Oxides form basic solutions in water:

MgO = Mg²⁺ + O^2-

O^2- + H₂O = 2OH^-

Covalent Oxides form acidic solutions in water:

SO³ + H₂O = H₂SO₄

H₂SO₄ = H⁺ + HSO₄^- 

B. Periodic trends in electron configurations, atomic properties such as radius, electronegativity, ionization potential, and chemical reactivity

Look at the Periodic Table, the vertical columns of the charts are called groups, while the rows are referred to as periods.

Metals, non-mental, and Atomic Radius

The elements in the periodic table can be classified into three groups based on their physical properties; metals, nonmetals and metalloids.

The metals are the largest group of elements. They are;

·         All solids (except mercury);

·         Have a lustrous (shiny) appearance;

·         Are malleable and ductile;

·         Good conductors of heat and electricity.

The nonmetals. The nonmetals;

·         Found in gas, liquid or solid state;

·         Lack the remaining properties of the metals, i.e. are not malleable or ductile and are generally poor conductors (graphite is a very good electrical conductor).

The metalloids are a small collection of elements that lie between the metals and the nonmetals in the periodic table and share some of the properties of metals and nonmetals.

Atomic size

Atomic Size: decreases going from left right and from bottom top.

• Size goes up with atomic number for any individual group.
• Size decreases irregularly as atomic number increases for any given period (more charge pulls electrons in to nucleus, but shielding reverses as subshells [s or p orbital sets] fill.