Elements of construct in biology and Paper Structure

The Periodic Table and Atomic Structure

Chemistry notes by Mr. Azabo Kennedy · Tel: 0786535915

Introduction Metals and non-metals Important groups Structure of an atom Atomic and mass number Electronic configuration Valence electrons Ions Summary

Introduction

Definition: The periodic table is the tabular arrangement of the chemical elements, based on their atomic number, electronic configuration and recurring chemical properties.

It provides a valuable tool for understanding the relationships between elements and predicting their behavior. The periodic table is made up of rows and columns.

• The horizontal rows are called periods. They are numbered according to the number of electron shells (energy levels) that are filled by electrons: 1, 2, 3, etc.
• The vertical columns are called groups. The group number is written in Roman numerals on top of the column and represents the number of electrons in the outermost shell (energy level).

Layout: metals sit on the left and in the centre, non-metals on the right, and the metalloids lie along the stair-step line between them.

Classes of elements

Elements in the periodic table are classified as metals, metalloids, non-metals and noble (inert or rare) gases.

Metals

Metals are on the left of the periodic table, in groups I, II and III. They react by losing electrons. Their properties:

• Shiny and reflective (lustrous).
• Good conductors of heat and electricity.
• High density.
• Tend to lose electrons to form positive ions (cations).
• React with oxygen and water to form oxides, leading to corrosion.

Examples of metals:

• Alkali metals: lithium, sodium, potassium.
• Alkaline earth metals: magnesium, calcium, barium.
• Transition metals: iron, copper, zinc.

Non-metals

Non-metals are on the right-hand side of the periodic table, in group IV to group VIII. Their properties:

• Dull and non-reflective.
• Non-ductile.
• Poor conductors of heat and electricity, except carbon in the form of graphite.
• React by gaining electrons to form negative ions (anions).

Examples of non-metals:

• Halogens: chlorine, iodine, fluorine, etc.
• Noble gases: not reactive because they are very stable. Helium, neon, argon, etc.
• Other non-metals: carbon, sulphur, phosphorus.

Transition elements

Found between group II and group III and towards the bottom of the periodic table. They form coloured compounds and have variable valences. Examples: copper, iron, chromium, cobalt.

Metalloids (semi-metals)

Found on the stair-step line that divides metals from non-metals. Their properties are intermediate between those of metals and non-metals. Important examples: silicon, germanium.

Important groups and their properties

Group	Name	Properties
I	Alkali metals	Highly reactive metals that readily lose one electron.
II	Alkaline earth metals	Reactive metals that lose two electrons to form ions.
VII	Halogens	Highly reactive non-metals that gain one electron to react.
VIII	Noble (inert) gases	Non-reactive gases with a full outer electron shell.

The structure of an atom

When scientists started exploring matter, they realized it can be divided into smaller and smaller particles. They called the smallest particle an atom, from the Greek word atomos, meaning "indivisible". They found that the atom keeps its chemical identity through all chemical and physical changes.

Definition: An atom is the smallest, indivisible and electrically neutral particle of an element that can take part in a chemical reaction.

Nature of the atom

The atom is the smallest particle of an element. It is made up of sub-atomic particles: electrons, protons and neutrons. Atoms of different elements differ because they have different numbers of sub-atomic particles.

The protons and neutrons are in the nucleus (centre) of the atom, and the electrons orbit outside it in shells (energy levels). Energy levels are counted from the centre outwards as 1st, 2nd, etc.

The atomic number and mass number

The nucleus is made up of protons and neutrons, which are together called nucleons.

Particle	Relative mass	Electric charge	Comments
Proton	1	+1 (positive)	In the nucleus (a nucleon)
Neutron	1	0 (zero)	In the nucleus (a nucleon)
Electron	1/1850	−1 (negative)	Arranged in energy levels (shells) around the nucleus

Since an atom is electrically neutral, the number of protons in the nucleus equals the number of electrons. This number is the atomic number, symbol Z. The atomic number is the number of protons in an atom.

The total number of protons and neutrons in one atom is its mass number (A).

Mass number (A) = number of protons (Z) + number of neutrons (n)
Mass number (A) = atomic number (Z) + number of neutrons (n)

Symbols: the mass (nucleon) number A is written as a superscript at the top-left of the atom's symbol. The atomic number Z is written as a subscript at the bottom-left. For carbon this is ¹²₆C.

Example

The symbol of a sodium atom with atomic mass 23 and atomic number 11 is Na. Calculate the number of (a) protons, (b) electrons and (c) neutrons.

Atomic number Z = 11, atomic mass A = 23

• (a) No. of protons = Z = 11
• (b) No. of electrons = 11
• (c) No. of neutrons = A − Z = 23 − 11 = 12

Electronic configuration

Electronic configuration is the arrangement of electrons in the shells (energy levels) around the nucleus of an atom. Since an atom is neutral, the number of electrons equals the atomic number (Z).

Rules for filling shells

• Electrons fill the shell closest to the nucleus first (the 1st shell), then move outwards: 1st, 2nd, 3rd, and so on.
• The 1st shell holds a maximum of 2 electrons (1 pair). The 2nd, 3rd, 4th, 5th, etc. hold a maximum of 8 electrons. If more electrons are still available, they occupy the next shells.

How to write it

1. Find the atomic number (Z), which is the number of electrons.
2. Fill the shells in order, following the maximum capacities above.
3. Write the numbers separated by commas, colons or semicolons, e.g. 2, 8, 1 or 2:8:1 or 2;8;1.

Example: sodium (Na), Z = 11

• 1st shell: 2 (full)
• 2nd shell: 8 (full)
• 3rd shell: 1 (half full)

Therefore the electronic configuration is 2:8:1.

Exercise: write the electronic configuration of the following atoms.

1. Helium
2. Hydrogen
3. Magnesium
4. Iodine
5. Chromium

Valence electrons and the periodic table

• Valence electrons are the electrons in the outermost shell. They determine how an element reacts.
• The number of occupied shells equals the period number.
• The number of valence electrons equals the group number (for groups 1 to 7). For example, sodium (2, 8, 1) is in Period 3, Group 1.
• Elements in the same group have the same number of valence electrons, so they have similar chemical properties.

Stable configurations

Atoms with a full outer shell are stable and unreactive. They normally have 8 electrons in the outermost shell. These are the noble gases: helium (2), neon (2, 8) and argon (2:8:8).

Other atoms are unstable. To become stable, they react by gaining or losing electrons to achieve a full outer shell, like the noble gases.

• Metals tend to lose electrons to form positive ions. Na (2:8:1) becomes Na⁺ (2:8) after losing one electron.
• Non-metals tend to gain electrons to form negative ions. Cl (2:8:7) becomes Cl⁻ (2:8:8) after gaining one electron.

Ions

Positive and negative ions form when electrons are transferred between atoms. This leaves an imbalance between the number of protons and electrons. Atoms lose or gain electrons to reach a stable electronic structure equal to that of the noble gases.

Whether an atom loses or gains electrons depends on its number of valence electrons. Atoms with fewer than 4 valence electrons find it easier to lose them in chemical reactions, while those with more than 4 find it easier to gain electrons. The charge on a cation (positive ion) shows the number of electrons donated; the charge on an anion (negative ion) shows the number gained.

Formation of positively charged ions

Atoms of metals form positive ions by losing 1 or more electrons from their valence (outermost) shell. The atom then has fewer electrons than protons, and because protons are positive, it has an overall positive charge. An ion with a positive charge is called a cation.

Sodium: Na → Na⁺   (11 protons, 11 electrons: 2, 8, 1 → 11 protons, 10 electrons: 2, 8)

Atoms, particularly metals, lose electrons to reach a stable configuration like that of a noble gas. Cations have empty valence shells and are smaller than their neutral atoms.

Formation of negatively charged ions

Atoms of non-metals form negative ions by gaining 1 or more electrons, which fills their outermost energy level. An ion with a negative charge is called an anion.

Chlorine: Cl → Cl⁻   (17 protons, 17 electrons: 2, 8, 7 → 17 protons, 18 electrons: 2, 8, 8)

Once they form ions, they attain the stable electronic configuration of noble gases such as helium, krypton and xenon.

Electronic configuration of ions

Exercise: write the electronic configuration of the following ions.

1. Sodium ion
2. Calcium
3. Potassium
4. Copper (I) ion
5. Chromium (III) ion
6. Lead (II) ion
7. Chlorine
8. Hydrogen
9. Carbon
10. Nitrogen

In summary

• Metals belong to groups such as group I, group II and group III. They gain stability by losing their outermost electrons during chemical reactions.
• Non-metals belong to groups such as group V, VI and VII. They gain stability by gaining electrons.
• Group VIII elements are stable and do not take part in chemical reactions.

Author: Mr. Azabo Kennedy · Tel: 0786535915

Converted from the original PDF notes. Diagrams are not included.