Arrangement of elements Revision Notes

Each element in the Periodic Table is arranged according to its atomic number, which is the number of protons in the nucleus of its atoms. The atomic number increases from left to right across each row, ensuring a unique position for every element.

This ordering by atomic number, rather than atomic mass, provides a consistent pattern of chemical properties and corrects earlier inconsistencies found in earlier tables arranged by atomic mass.

• Atomic number = number of protons in an atom's nucleus.
• Elements are placed in order of increasing atomic number.
• This order determines the element's position in the table.

The Periodic Table is organised into horizontal rows called periods. Each period corresponds to the filling of a new electron shell around the nucleus.

As you move across a period from left to right, the atomic number increases by one for each element, and electrons are added to the same outer shell.

The properties of elements gradually change across a period due to the increasing nuclear charge affecting electron behaviour.

• Periods are horizontal rows in the Periodic Table.
• Each period represents a principal energy level (electron shell).
• Atomic number increases by one across a period.

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 properties depend largely on valence electrons, elements in the same group exhibit similar chemical behaviour, such as forming ions with the same charge or reacting similarly with other substances.

Groups are numbered from 1 to 18 in the modern IUPAC system, but older systems use group numbers 1 to 8 with A and B subdivisions.

• Groups are vertical columns in the Periodic Table.
• Elements in a group have the same number of valence electrons.
• Elements in the same group have similar chemical properties.

The arrangement by atomic number leads to a periodic pattern of chemical properties, known as the periodic law.

Elements with similar valence electron configurations recur at regular intervals, which is why elements in the same group behave similarly.

For example, all Group 1 elements have one valence electron and are highly reactive metals, while Group 17 elements have seven valence electrons and are reactive non-metals.

• Periodic law: properties repeat periodically when elements are arranged by atomic number.
• Valence electron number determines chemical properties.
• Groups reflect recurring chemical behaviour.

Sodium (Na) has atomic number 11, meaning it has 11 protons and 11 electrons. Its electron configuration is 2,8,1, so it has one valence electron.

Chlorine (Cl) has atomic number 17, with 17 protons and electrons. Its electron configuration is 2,8,7, so it has seven valence electrons.

In the Periodic Table, sodium is placed in Period 3 (third row) and Group 1 (first column), reflecting its single valence electron in the third shell. Chlorine is also in Period 3 but in Group 17, reflecting its seven valence electrons in the same shell.

• Sodium: atomic number 11, electron configuration 2,8,1, Period 3, Group 1.
• Chlorine: atomic number 17, electron configuration 2,8,7, Period 3, Group 17.
• Position reflects increasing atomic number and valence electrons.

Elements in the Periodic Table are arranged in periods (rows) and groups (columns) according to increasing atomic number. Across a period, elements on the left are metals, while those on the right are non-metals. The transition from metallic to non-metallic properties occurs gradually as you move from left to right across a period. This change reflects differences in atomic structure and electron arrangements that influence physical and chemical properties.

• Metals occupy the left side of a period.
• Non-metals are found on the right side of a period.
• Elements in the middle of a period may show metalloid properties.

Metallic elements typically have high electrical and thermal conductivity, are malleable and ductile, and have a shiny appearance. Non-metals generally have low electrical conductivity, are brittle in solid form, and lack metallic lustre. As you move across a period, electrical conductivity decreases, and malleability is lost, reflecting the change from metallic to non-metallic character.

• Metals: good conductors, malleable, ductile, lustrous.
• Non-metals: poor conductors, brittle, dull appearance.
• Conductivity decreases from left to right across a period.

Across a period, the number of protons in the nucleus (nuclear charge) increases, while electrons are added to the same principal energy level. The increasing nuclear charge pulls electrons closer, reducing atomic radius. This stronger attraction makes it harder for atoms to lose electrons, decreasing metallic character. Instead, atoms tend to gain or share electrons, increasing non-metallic character.

• Nuclear charge increases across a period.
• Atomic radius decreases due to stronger attraction.
• Metals lose electrons easily; non-metals gain or share electrons.

Metals tend to lose electrons to form positive ions (cations). For example, sodium (Na) loses one electron to form Na⁺:

Na → Na⁺ + e⁻ (oxidation) Non-metals tend to gain electrons to form negative ions (anions). For example, chlorine (Cl) gains one electron to form Cl⁻:

Cl + e⁻ → Cl⁻ (reduction) This difference in electron transfer behaviour reflects the transition from metallic to non-metallic properties across a period.

• Metals form positive ions by electron loss.
• Non-metals form negative ions by electron gain.
• Oxidation is electron loss; reduction is electron gain.

Sodium (Na), a metal on the left of period 3, loses one electron to form Na⁺. Chlorine (Cl), a non-metal on the right of period 3, gains one electron to form Cl⁻. When these ions combine, they form sodium chloride (NaCl) via ionic bonding:

Na → Na⁺ + e⁻ (oxidation) Cl + e⁻ → Cl⁻ (reduction) Na⁺ + Cl⁻ → NaCl (ionic compound) This illustrates the contrasting metallic and non-metallic behaviour across a period.

• Sodium loses one electron to form Na⁺.
• Chlorine gains one electron to form Cl⁻.
• Ionic bond forms between Na⁺ and Cl⁻ to make NaCl.

Elements in the periodic table are arranged in groups (vertical columns) that reflect the number of electrons in their outermost shell, called valence electrons.

The group number for main group elements (Groups 1, 2, and 13 to 18) corresponds to the number of valence electrons an atom has.

For example, Group 1 elements have 1 valence electron, Group 2 have 2, Group 13 have 3, and so on up to Group 18, which have 8 valence electrons (except helium with 2).

• Valence electrons determine chemical reactivity and ion formation.
• Atoms tend to gain or lose electrons to achieve a stable electronic configuration, usually a full outer shell (octet).

Atoms form ions by losing or gaining electrons to reach a stable electron arrangement similar to the nearest noble gas.

Metals, typically found in Groups 1, 2, and 13, lose electrons to form positive ions (cations).

Non-metals, typically in Groups 15, 16, and 17, gain electrons to form negative ions (anions).

• Loss of electrons results in a positive charge equal to the number of electrons lost.
• Gain of electrons results in a negative charge equal to the number of electrons gained.
• For example, sodium (Group 1) loses 1 electron to form Na⁺.
• Chlorine (Group 17) gains 1 electron to form Cl⁻.

The charge of the ion formed by an element can be predicted from its group number by considering whether it loses or gains electrons to complete its outer shell.

For metals (Groups 1, 2, and 13), the ion charge is positive and equals the group number.

For non-metals (Groups 15, 16, and 17), the ion charge is negative and equals 8 minus the group number, representing the electrons gained to complete the octet.

• Group 1 elements form +1 ions (e.g., Na⁺).
• Group 2 elements form +2 ions (e.g., Mg²⁺).
• Group 13 elements form +3 ions (e.g., Al³⁺).
• Group 15 elements form -3 ions (e.g., N³⁻).
• Group 16 elements form -2 ions (e.g., O²⁻).
• Group 17 elements form -1 ions (e.g., Cl⁻).

Magnesium is in Group 2 and has 2 valence electrons.

To achieve a stable electronic configuration, magnesium loses 2 electrons.

This forms a Mg²⁺ ion with a full outer shell like neon.

• Equation for ion formation: Mg (s) → Mg²⁺ (aq) + 2e⁻
• Magnesium ion charge is +2, matching its group number.
• This loss of electrons is oxidation.

When metals react with non-metals, electrons transfer to form ions with charges predicted by their group numbers.

Example: Sodium reacts with chlorine to form sodium chloride.

The balanced ionic equation is: 2Na (s) + Cl₂ (g) → 2Na⁺ (aq) + 2Cl⁻ (aq) → 2NaCl (s)

• Sodium loses 1 electron per atom forming Na⁺ ions.
• Chlorine gains 1 electron per atom forming Cl⁻ ions.
• The ionic bond is the electrostatic attraction between Na⁺ and Cl⁻ ions.
• Observation: Sodium metal reacts vigorously with chlorine gas, producing white crystalline sodium chloride.

The chemical properties of elements are primarily determined by their electron arrangements, especially the number of electrons in the outermost shell, called valence electrons. Elements in the same group of the Periodic Table have the same number of valence electrons, which results in similar chemical behaviour.

For example, all elements in Group 1 have one electron in their outer shell, while all elements in Group 17 have seven valence electrons. This commonality explains why elements in a group tend to react in similar ways and form ions with the same charge.

• Valence electrons determine chemical reactivity.
• Elements in the same group have identical valence electron counts.
• Similar valence electron configurations cause similar chemical properties.

Group 1 elements (lithium, sodium, potassium, rubidium, caesium, and francium) are metals with one valence electron. They are highly reactive, especially with water, because they tend to lose this electron to form +1 ions (M⁺).

When alkali metals react with water, they form metal hydroxides and hydrogen gas. The reaction is vigorous and exothermic, often producing enough heat to ignite the hydrogen.

Example reaction with sodium: 2Na(s) + 2H₂O(l) → 2NaOH(aq) + H₂(g)

• Group 1 metals have one valence electron.
• They lose one electron to form M⁺ ions.
• React vigorously with water to form hydroxides and hydrogen.
• Reactivity increases down the group due to easier electron loss.

Group 17 elements (fluorine, chlorine, bromine, iodine, and astatine) are nonmetals with seven valence electrons. They tend to gain one electron to complete their outer shell, forming -1 ions (X⁻).

Halogens react with metals to form ionic salts, where the halogen ion carries a negative charge.

Example reaction with chlorine and sodium: 2Na(s) + Cl₂(g) → 2NaCl(s)

• Group 17 elements have seven valence electrons.
• They gain one electron to form X⁻ ions.
• React with metals to form ionic salts.
• Reactivity decreases down the group as atoms gain electrons less easily.

Because elements in the same group have the same number of valence electrons, they form ions with the same charge. Group 1 elements form +1 ions, Group 2 form +2 ions, and Group 17 form -1 ions.

This pattern helps predict the formulae of compounds formed by these elements. For example, sodium (Na⁺) combines with chloride (Cl⁻) to form NaCl.

Reactivity trends within groups are influenced by atomic size and electron shielding. For metals, reactivity increases down the group as valence electrons are lost more easily. For nonmetals, reactivity decreases down the group as gaining electrons becomes harder.

• Elements in a group form ions with the same charge.
• Ionic compound formulae can be predicted from ion charges.
• Reactivity trends vary down groups due to atomic structure changes.

Potassium (K) is a Group 1 metal with one valence electron. When it reacts with water, it loses this electron to form K⁺ ions, producing potassium hydroxide and hydrogen gas.

Balanced equation: 2K(s) + 2H₂O(l) → 2KOH(aq) + H₂(g)

Observations include effervescence (bubbling), the metal melting due to heat released, and the solution becoming alkaline (pH > 7).

• Potassium loses one electron to form K⁺.
• Reaction with water produces KOH and H₂.
• Reaction is vigorous and exothermic.
• Solution turns alkaline due to KOH formation.

The Periodic Table arranges elements in order of increasing atomic number (number of protons). Elements are placed in rows called periods and columns called groups.

Each group contains elements with the same number of electrons in their outermost shell (valence electrons), which largely determines their chemical behaviour.

Periods correspond to the filling of electron shells; as you move across a period, the number of protons and electrons increases by one for each successive element.

• Groups are vertical columns with elements having similar chemical properties.
• Periods are horizontal rows showing gradual changes in properties.
• The group number (for main-group elements) equals the number of valence electrons.

Elements in the same group have the same number of valence electrons, resulting in similar chemical properties and bonding behaviour.

For example, Group 1 elements (alkali metals) have one valence electron, making them highly reactive metals that form +1 ions by losing that electron.

Group 17 elements (halogens) have seven valence electrons and tend to gain one electron to form -1 ions, making them reactive non-metals.

• Group 1 elements react vigorously with water to form hydroxides and hydrogen gas: 2Na (s) + 2H₂O (l) → 2NaOH (aq) + H₂ (g)
• Group 17 elements react with metals to form ionic halide salts, e.g. 2Na (s) + Cl₂ (g) → 2NaCl (s)

Reactivity in groups changes predictably due to atomic structure changes down the group.

In Group 1, reactivity increases down the group because the outer electron is further from the nucleus and more easily lost.

In Group 17, reactivity decreases down the group because the atoms are larger and less able to attract an extra electron.

• Lithium reacts less vigorously with water than potassium.
• Chlorine is more reactive than iodine in displacement reactions.

Moving across a period from left to right, elements change from metals to non-metals.

Atomic radius decreases across a period due to increasing nuclear charge pulling electrons closer.

Ionisation energy generally increases across a period, making atoms less likely to lose electrons and more likely to gain or share electrons.

• Sodium (Group 1) is a metal with low ionisation energy.
• Chlorine (Group 17) is a non-metal with high ionisation energy.

Electron configuration shows the distribution of electrons in shells and subshells.

Elements in the same group have similar outer electron configurations, explaining their similar chemical properties.

For example, all Group 2 elements have the configuration ending in s², making them reactive metals forming +2 ions.

• Calcium: [Ar] 4s², forms Ca²⁺ ions.
• Oxygen: [He] 2s² 2p⁴, tends to gain 2 electrons to form O²⁻ ions.

The Periodic Table arranges elements in vertical columns called groups. Elements within the same group have the same number of electrons in their outermost shell, known as valence electrons. This similarity in valence electron configuration results in elements exhibiting comparable chemical properties.

For example, Group 1 elements (alkali metals) all have one valence electron, making them highly reactive metals that form similar compounds with nonmetals. Recognising these patterns is essential for predicting element behaviour and reactions.

• Groups are vertical columns in the Periodic Table.
• Elements in a group share the same number of valence electrons.
• Valence electrons determine chemical properties and reactivity.

The valence electrons are primarily responsible for an element's chemical behaviour because they participate in bond formation. Elements in the same group have identical valence electron configurations, which explains their similar chemical properties.

For instance, Group 7 elements (halogens) have seven valence electrons and tend to gain one electron to form -1 ions, resulting in similar chemical reactions such as forming halide salts with metals.

The general electron configuration for Group 1 elements is ns¹, for Group 2 is ns², and for Group 7 is ns² np⁵, where n represents the period number.

• Valence electrons are involved in bonding and reactions.
• Same valence electron configuration leads to similar chemical properties.
• Example: Group 1 elements have configuration ns¹ and are highly reactive metals.
• Example: Group 7 elements have configuration ns² np⁵ and form -1 ions.

Reactivity within groups often follows a predictable trend as you move down the group. This is due to changes in atomic size and shielding effect, which affect how easily atoms lose or gain electrons.

In Group 1, reactivity increases down the group because the outer electron is further from the nucleus and more easily lost, forming positive ions.

In Group 7, reactivity decreases down the group because the atoms find it harder to gain an extra electron due to increased atomic radius and shielding.

• Reactivity trends depend on atomic size and electron shielding.
• Group 1 reactivity increases down the group (easier electron loss).
• Group 7 reactivity decreases down the group (harder electron gain).

Physical properties such as melting point, boiling point, and density also show patterns within groups. For example, in Group 1, melting and boiling points decrease down the group, while density generally increases.

These trends arise from changes in atomic structure and bonding forces as atomic size increases.

Recognising these physical property patterns helps in identifying elements and predicting their states and uses.

• Melting and boiling points vary predictably within groups.
• Group 1 melting and boiling points decrease down the group.
• Density tends to increase down the group due to heavier atoms.

Given data for Group 1 elements: lithium, sodium, potassium, rubidium, and caesium, showing melting points and reactivity with water, patterns can be identified.

For example, lithium reacts slowly with cold water, producing lithium hydroxide and hydrogen gas:

2Li(s) + 2H₂O(l) → 2LiOH(aq) + H₂(g) ↑ (balanced equation) Observations: fizzing, hydrogen gas evolution, solution becoming alkaline (tested with universal indicator). As you move down the group, the reaction becomes more vigorous, with potassium reacting explosively with cold water. This pattern reflects increasing reactivity down Group 1 due to easier electron loss.

• Balanced equation for lithium and water reaction: 2Li(s) + 2H₂O(l) → 2LiOH(aq) + H₂(g)
• Observations: fizzing, hydrogen gas, alkaline solution.
• Reactivity increases down Group 1, shown by reaction vigour.