The Periodic Table

• The first periodic table was created by Dmitri Mendeleev in 1869, containing 69 elements.
• The modern periodic table contains 115 elements.
• Mendeleev arranged elements by increasing atomic mass, while the modern table arranges them by increasing atomic number (number of protons).

Periods and Groups

• Periods are horizontal rows of elements.
• Groups are vertical columns of elements, numbered from I to VIII or 0.
• Elements between Groups II and III are called transition metals.

Patterns in the Periodic Table

• Elements in the same group have the same number of valence electrons, matching the group number (e.g., Group II elements have 2 valence electrons).
• The charges on ions correspond to the number of valence electrons and group number.
• Elements on the left (Groups I, II, III) lose electrons to form cations.
• Elements on the right (Groups V, VI, VII) gain electrons to form anions.
• Group IV elements can either lose or gain electrons depending on the reacting element.
• Transition metals can form multiple cations (e.g., 2+, 3+).

Bonding

Elements in the same group form the same type and number of bonds because they have the same number of valence electrons. For example, sodium (Group I) forms NaCl, and other Group I elements form similar compounds (RbCl, KCl, LiCl, CsCl).

Metals and Non-metals

• Elements to the left of Group IV are metals, while those to the right are non-metals.
• Group IV elements are metalloids, exhibiting both metallic and non-metallic properties.

Changes within Groups

• The atomic (proton) number increases as you move down a group.
• The rate of increase in proton number is gradual for elements on either side of the table but more significant among transition metals.

Using the Periodic Table to Predict Properties

• Formulas and structures of compounds can be predicted.
• For example, since chlorine (Cl₂), iodine (I₂), and bromine (Br₂) are diatomic molecules in Group VII, fluorine would also form F₂.

Group I Elements (The Alkali Metals)

• React with water to form alkaline solutions.
• Turn red litmus paper blue.
• Are the most reactive metals.
• Have one valence electron.
• Are shiny, silvery solids that are soft and easily cut with a knife.
• Have low densities and low melting points, which increase down the group.
• React quickly with air and are stored in oil.
• React vigorously (sometimes explosively) with cold water.
• Form ionic compounds with a +1 charge.
• Have similar chemical formulas.
• Become more reactive down the group.

Group VII Elements (The Halogens)

• React with most metals to form salts.
• Are very reactive.
• Have seven valence electrons.
• Form diatomic molecules (e.g., F₂).
• Become darker and more solid moving down the group.
• Have low melting and boiling points, which increase down the group.
• Are poisonous.

Compounds of the Halogens

• Halogens form -1 charged ions and similar formulas, e.g., NaBr, NaI.
• They react vigorously with metals to form ionic salts.
• Example reaction:

  $$ 2K + Br_2 \rightarrow 2KBr $$

• Reactivity decreases down the group.

Displacement Reactions

A more reactive halogen can displace a less reactive halogen from its salt solution. For example:

$$ F_2(aq) + 2NaBr(aq) \rightarrow 2NaF(aq) + Br_2(aq) $$

Group VIII or 0 Elements (The Noble Gases)

• Are the least reactive elements.
• Do not easily form bonds.
• Have complete outer electron shells, making them very stable.
• Exist as monatomic gases (single atoms).
• Have low melting and boiling points.

Uses of the Noble Gases

• Argon is used in light bulbs to prevent the filament from reacting with air.
• Neon is used in advertising lights.
• Helium is used in balloons and airships because it is lighter than air.

Properties of Transition Elements

• All first transition series elements are metals.
• They have high melting points and densities.
• They exhibit variable oxidation states, e.g., Fe²⁺ and Fe³⁺.
• They form colored compounds, e.g., CuSO₄ (blue) and FeSO₄ (green).
• They form complex ions, e.g., MnO₄^-.
• They act as catalysts in chemical reactions.

Uses of Transition Elements

• Many transition metals and their compounds are catalysts.
• Iron is used in the Haber process to manufacture ammonia.
• Vanadium(V) oxide is used in the Contact process to produce sulfuric acid.
• Nickel is used in the hydrogenation of alkenes to make margarine.

Advantages of Using Transition Elements

• They speed up chemical reactions, saving time in industrial processes.
• They reduce energy consumption, lowering production costs and conserving energy resources like oil.

Variation of Atomic Properties Across a Period and Down a Group

Atomic properties, such as atomic radius, change as you move across a period or down a group in the periodic table. Since electrons are distributed in a cloud around the nucleus with no clear boundary, the exact size of an atom is hard to define.

1. Atomic Radius

Atomic radius is the distance between the nuclei of two identical atoms bonded together. There are two main types:

• Covalent Radius: Half the distance between two identical bonded atoms.
• Van der Waals Radius: Half the distance between two identical non-bonded atoms.

Trends in atomic radius:

• Atomic radius increases down a group because new electron shells are added, making atoms larger even though nuclear charge increases.
• Atomic radius decreases across a period because increasing nuclear charge pulls electrons closer to the nucleus.

Examples: Potassium has a larger radius than sodium; caesium is larger than rubidium.

2. Ionic Radius

Ions form when atoms gain or lose electrons:

• Cations (positive ions) are smaller than their parent atoms because the loss of electrons increases the effective nuclear charge.
• Anions (negative ions) are larger than their parent atoms because gaining electrons reduces the effective nuclear charge.

Across the second period:

• Cationic radii decrease from sodium to aluminum.
• Anionic radii increase from phosphorus to chlorine.

3. Ionization Energy

Ionization energy is the energy needed to remove an electron from a gaseous atom. It is measured in kJ/mol.

Trends in ionization energy:

• It increases across a period because the nuclear charge increases, pulling electrons closer to the nucleus.
• It decreases down a group because outer electrons are farther from the nucleus and are more shielded by inner electrons.

Factors affecting ionization energy:

• Distance of outer electrons from the nucleus.
• Effective nuclear charge.
• Screening effect by inner electrons.

4. Electron Affinity

Electron affinity is the energy released when a gaseous atom accepts an electron, expressed in kJ/mol or electron volts (eV).

Trends in electron affinity:

• It increases across a period from left to right.
• It decreases down a group.

Group 1 elements (alkali metals) have the lowest electron affinity, while Groups VI and VII elements have the highest. Noble gases have full electron shells and do not readily accept electrons.

5. Electronegativity

Electronegativity is the ability of an atom in a molecule to attract shared electrons. It is most notable in molecules made of different elements.

Trends in electronegativity:

• It increases across a period because of increasing nuclear charge and decreasing atomic size.
• It decreases down a group due to increased atomic size and greater shielding effect.

Fluorine has the highest electronegativity. Noble gases generally do not have electronegativity values because their outer electron shells are complete.