CHEMISTRY FORM 2
- 1.1 Structure of the atom
- 1.2 Atomic Number and Mass Number
- 1.3 Isotopes
- 1.4 Energy levels and electron arrangement
- 1.5 Development of the Periodic Table
- 1.6 Relative Atomic Mass and Isotopes
- 1.7 Ion Formation
- 1.8 Chemical Formulae
- 1.9 Chemical Equations
- 2.1 Alkali metals (Group I elements)
- 2.2 Alkali Earth Metals (Group II elements)
- 2.3 Halogens (Group VII elements)
- 2.4 Noble gases (Group VIII elements)
- 2.5 Properties and Trends Across the Periodic Table
- 3.1 Bond
- 3.2 Ionic bond
- 3.3 Giant ionic structure
- 3.4 Covalent bond
- 3.5 Co-ordinate bond
- 3.6 Molecular structures
- 3.7 Giant covalent structures
- 3.8 Metallic Bond
- 3.9 Types of bond across a period
- 3.10 Oxides of elements in Period 3
- 3.11 Chlorides of Period 3 elements
- 4.1 What is a salt?
- 4.2 Types of salt
- 4.3 Solubility of salts in water
- 4.4 Methods of preparing salts
- 4.4.1 Reacting a Metal with an Acid
- 4.4.2 Reacting an Acid with a Base (Neutralization)
- 4.4.3 Reacting an Acid with a Carbonate (or hydrogencarbonate of metal)
- 4.4.4 Combining elements Directly (Direct Combination of elements)
- 4.4.5 Precipitation (Double decomposition)
- 4.5 Action of heat on salts
- 4.6 Uses of salts
- 5.1 Electrical conduction
- 5.2 Electrical conductivity of molten substances
- 5.3 Electrical conductivity of substances in aqueous state
- 5.4 Electrolysis
- 5.5 Applications of electrolysis
- 6.1 Allotropes of carbon
- 6.2 Chemical properties of carbon
- 6.3 Carbon (IV) oxide
- 6.4 Carbon (II) oxide (CO)
- 6.5 Large scale production of sodium carbonate and sodium hydrogencarbonate
- 6.6 Effect of carbon (II) oxide and carbon (IV) oxide on the environment
- 6.7 Carbon cycle
Structure of the Atom, and the Periodic Table: Relative Atomic Mass and Isotopes
1.0 Structure of the Atom, and the Periodic Table
1.7 Ion Formation
Analogy: Behavior is caused. It aims to satisfy a need. We drink water to quench our thirst, eat when hungry and so on. Similarly, atoms react to complete their outermost energy level and be stable. Elements with complete outermost energy levels are therefore unreactive (because they are stable).
Two of the ways by which atoms can become stable are as follows.
- Losing all the outermost electrons if they are few (1 to 3) so that they remain with the complete inner energy level.
- Gaining electrons (1 to 3) if the energy level is nearly complete, with 7, 6, or 5 electrons respectively.
This is because losing 1 to 3 electrons requires less energy than gaining the balance (7 to 5) to reach 8 electrons. Also, gaining 1 to 3 electrons requires less energy than losing 5 to 7. When an atom gains or loses electrons, the product is electrically charged and is called an ion.
An ion is an electrically charged particle formed when an atom loses or gains electrons.
An ion formed by losing an electron is positively charged (+) and is called a cation. One that is formed by gaining an electron is negatively charged (-) and is called an anion. Watch the pencast on ion formation.
(courtesy Youtube-Ion formation OK by Joseph Rabari)
Refer to the Periodic Table as you answer the following questions
- Identify 3 unreactive elements among the first 20 elements (A = 1 to 20). In which group are these elements found?
- Name all the elements within the atomic numbers 1 to 20, which are expected to react by losing electrons. Identify the groups to which these elements belong.
- Name all the elements within the atomic numbers 1 to 20, which are expected to react by losing electrons. What are the groups to which these elements belong?
Answers to Questions 1.7(a)
An atom is electrically neutral. But when it loses an electron (negatively charged), the total number of electrons becomes less than the protons (positively charged) in the nucleus. The product therefore has less negative than positive charges. It is a positively charged ion. We can represent this change generally as
X − e ⟶ X+
This equation reads, "atom X loses (-) an electron (e) to form a positive ion, X+."
Example: Li − e ⟶ Li+
NB: We can as well move the electron (e) to the right of the equation, in which case we add rather than subtract it (as is the case when crossing equal sign). That is,
Li ⟶ Li+ + e
Positive ions (cations), are formed by metals (Group I to III elements) because they react by losing electrons.
When an atom gains an electron, the total number of electrons (-) becomes more than the protons (+). The product therefore has more negative than positive charges. It is negatively charged (anion). We can represent this change generally as
Y + e ⟶ Y-
Example: F + e ⟶ F-
Negative ions are formed by non-metals (Group V to VII elements) because they react by gaining electrons.
- Indicate the number of electrons an atom of each of the following metals loses to form an ion. (a) sodium (b) magnesium (c) aluminium
- Write equations to represent ion formation by
- Sodium metal
- Magnesium metal
- Write equations to represent ion formation by
Answers to Questions 1.7(b)
Ion formation can also be represented on energy level diagrams. The following diagram shows ion formation by lithium atom.
Figure 1.7: Ion formation by lithium atom
Note that the outermost electron and energy level have been lost during ion formation. The resulting structure has a complete outermost energy level and therefore stable. It does not react further. Electronic arrangement changes from 2:1 to 2, and the ion is smaller.
Draw an energy level model to represent an ion of
- Fluorine, F (9p and 10n).
- Oxygen (8p and 8n)
- Sodium (11p and 12n)
- Magnesium (12p and 12n)
Answers to Questions 1.7(c)
Oxidation number (or oxidation state)
This is the state of charge on an ion. Positive and negative ions have positive and negative oxidation numbers respectively. Ions of sodium (Na+), magnesium (Mg2+),
aluminum (Al3+), chlorine (Cl-), and oxygen (O2-), for example, have oxidation numbers +1, +2, +3, -1, and -2 respectively. The sign comes before the number. For elements
and atoms in uncombined state, the oxidation number is zero (0).
NB: An ion has a similar electronic configuration as the noble gas atom in its period. For example, sodium ion and neon atom have the same electronic configuration 2:8. But they differ chemically because their nuclei have different numbers of protons; sodium has 11 and neon 10. It is the number of protons, and not electrons that identifies an element.