Arrangement of elements Revision Notes

The Periodic Table arranges elements in order of increasing atomic number, which is the number of protons in an atom's nucleus. This number uniquely identifies each element and determines its chemical identity. For example, hydrogen has atomic number 1, helium 2, lithium 3, and so on. Arranging elements by atomic number rather than atomic mass resolves inconsistencies seen in earlier tables, ensuring a consistent pattern of properties.

Because atomic number increases by one for each successive element, the table progresses from the simplest atoms to more complex ones. This order reflects the filling of electron shells around the nucleus, which influences chemical behaviour.

• Atomic number = number of protons in nucleus.
• Elements arranged left to right by increasing atomic number.
• Atomic number determines element identity and properties.

The Periodic Table is divided into horizontal rows called periods. Each period corresponds to the number of electron shells occupied by electrons in the atoms of the elements in that row.

For example, elements in period 1 have electrons in only the first shell, period 2 elements have electrons in two shells, and so forth. As you move across a period from left to right, electrons fill the outermost shell progressively.

This explains why elements in the same period show a gradual change in properties, such as increasing electronegativity or ionisation energy, as the outer shell fills.

• Periods run horizontally across the table.
• Period number = number of electron shells in atoms.
• Properties change gradually across a period as outer shell fills.

Vertical columns in the Periodic Table are called groups. Elements in the same group have the same number of electrons in their outermost shell, known as valence electrons.

Because chemical reactions mainly involve these outer electrons, elements in the same group exhibit similar chemical properties. For example, group 1 elements (alkali metals) all have one electron in their outer shell and react vigorously with water.

This similarity allows prediction of element behaviour based on group membership.

• Groups run vertically down the table.
• Elements in a group have the same number of valence electrons.
• Similar valence electron configuration leads to similar chemical properties.

Consider sodium (Na) and potassium (K). Sodium is in period 3, group 1, and potassium is in period 4, group 1. Both have one electron in their outer shell, so they react similarly, such as forming +1 ions.

In contrast, neon (Ne) and argon (Ar) are in group 0 (noble gases) but different periods (2 and 3). Both have full outer shells, making them chemically inert, but argon has an additional electron shell compared to neon.

These examples show how periods indicate shell number and groups indicate valence electron count, linking atomic structure to chemical properties.

• Sodium (period 3, group 1) and potassium (period 4, group 1) both have 1 valence electron.
• Neon (period 2, group 0) and argon (period 3, group 0) have full outer shells but different shell numbers.
• Period number = electron shells; group number = valence electrons.

By organising elements into periods and groups based on atomic number, the Periodic Table reveals periodic trends in element properties such as reactivity, atomic radius, and ionisation energy.

This arrangement helps chemists predict the behaviour of elements and their compounds, even for newly discovered or synthetic elements.

Understanding the relationship between atomic number, electron configuration, and element properties is fundamental to studying chemistry.

• Periodic Table shows repeating patterns (periodicity) in element properties.
• Arrangement by atomic number links atomic structure to chemical behaviour.
• Enables prediction of element properties and reactions.

In the periodic table, elements are arranged in periods (rows) according to increasing atomic number. Moving from left to right across a period, elements exhibit a gradual change in properties, transitioning from metallic to non-metallic character. This trend reflects changes in atomic structure, particularly the number of protons and electrons, which influence how atoms interact with other atoms and their environment.

• Elements on the left are metals; on the right, non-metals.
• Atomic number increases by one for each successive element.
• Properties change gradually, not abruptly, across a period.

As you move across a period, each element has one more proton in its nucleus and one more electron in its outer shell than the previous element. The increasing positive charge of the nucleus (nuclear charge) pulls the electrons closer, reducing atomic radius. Although electrons are added to the same shell, the stronger nuclear attraction is not fully offset by electron shielding, so the effective nuclear charge felt by outer electrons increases.

This stronger attraction affects how easily atoms lose or gain electrons, which is fundamental to their metallic or non-metallic behaviour.

• Nuclear charge increases by +1 per element across a period.
• Atomic radius decreases due to stronger attraction between nucleus and electrons.
• Electron shielding remains approximately constant within a period.

Elements on the left side of a period are metals. They tend to have few electrons in their outer shell (usually 1 to 3), which they can lose easily to form positive ions. This ability to lose electrons corresponds to metallic character. Metals are typically good conductors of heat and electricity because their outer electrons are delocalised and free to move.

They are also malleable and ductile due to metallic bonding, which allows layers of atoms to slide over each other without breaking bonds.

• Metals have low ionisation energies, facilitating electron loss.
• Good conductors of heat and electricity.
• Malleable and ductile due to metallic bonding.
• Example: Sodium (Na) loses one electron to form Na⁺.

Elements on the right side of a period are non-metals. They have more electrons in their outer shell (usually 5 to 8) and tend to gain or share electrons to complete their outer shell, forming negative ions or covalent bonds. Non-metals are poor conductors of heat and electricity because their electrons are held tightly and not free to move.

Physically, non-metals can be gases, liquids, or brittle solids at room temperature, unlike metals which are mostly solid and malleable.

• Non-metals have high ionisation energies and high electronegativities.
• Poor conductors of heat and electricity.
• Exist in various states: gases (e.g., oxygen), liquids (e.g., bromine), or brittle solids (e.g., sulfur).
• Example: Chlorine (Cl) gains one electron to form Cl⁻.

The transition from metallic to non-metallic properties across a period is caused primarily by the increasing nuclear charge, which pulls electrons closer and increases the energy required to remove electrons. This results in decreasing metallic character and increasing non-metallic character.

This trend explains why elements on the left lose electrons easily and those on the right tend to gain or share electrons. Understanding this helps predict element behaviour in chemical reactions and their physical properties.

• Increasing nuclear charge reduces atomic radius and increases ionisation energy.
• Metallic character decreases; non-metallic character increases across a period.
• Chemical reactivity changes accordingly: metals tend to form cations, non-metals form anions or covalent bonds.

The periodic table arranges elements into vertical columns called groups. Elements in the same group have similar chemical properties because they have the same number of electrons in their outermost shell, known as valence electrons. This similarity in electron configuration leads to comparable reactivity and bonding behaviour.

For example, Group 1 elements (alkali metals) all have one valence electron, making them highly reactive metals that readily lose this electron to form positive ions. In contrast, Group 18 elements (noble gases) have full outer shells, making them very unreactive or inert under normal conditions.

• Groups are vertical columns in the periodic table.
• Elements in a group share the same number of valence electrons.
• Similar valence electron configuration causes similar chemical properties.

Reactivity within a group can show clear trends. For instance, in Group 1, reactivity increases down the group because the outer electron is further from the nucleus and more easily lost. Conversely, in Group 7 (halogens), reactivity decreases down the group because atoms find it harder to gain an electron as atomic size increases.

Recognising these trends allows prediction of how vigorously an element will react. For example, potassium (K) is more reactive than sodium (Na) because its valence electron is less tightly held.

• Reactivity trends occur within groups.
• In Group 1, reactivity increases down the group.
• In Group 7, reactivity decreases down the group.
• Atomic size and electron shielding influence these trends.

The electron configuration of an element determines its chemical properties. Elements in the same group have identical numbers of electrons in their outer shell, which governs how they interact with other atoms.

For example, all Group 2 elements have two valence electrons and tend to form +2 ions by losing these electrons. This explains their similar reactions with water and acids, producing hydrogen gas and metal hydroxides or salts.

• Valence electrons determine chemical reactivity.
• Same group elements have identical valence electron counts.
• Ion formation patterns are similar within groups.

By recognising patterns in groups, chemists can predict how unknown or less familiar elements might behave. For example, if an element is in Group 1, it is expected to be a reactive metal that forms +1 ions.

This predictive power is useful in chemical synthesis and understanding reactions. For instance, knowing that chlorine (Group 7) is a reactive non-metal helps anticipate its reactions with metals to form salts.

• Group patterns enable prediction of element properties.
• Chemical reactions can be anticipated based on group membership.
• This aids in understanding and planning chemical experiments.