Atomic structure and the Periodic Table Revision Notes

At the centre of every atom lies the nucleus, a dense region containing two types of subatomic particles: protons and neutrons. Protons carry a positive electrical charge (+1), while neutrons are electrically neutral, carrying no charge. Together, these particles account for nearly all the atom's mass because they are much heavier than electrons.

The number of protons in the nucleus defines the element and is called the atomic number. For example, all carbon atoms have six protons. Neutrons contribute to the mass but do not affect the chemical identity. The total number of protons and neutrons is called the mass number. For instance, a carbon atom with six protons and six neutrons has a mass number of 12.

• Nucleus contains protons (+ charge) and neutrons (no charge).
• Protons determine the element's identity (atomic number).
• Neutrons add mass but no charge.
• Mass number = protons + neutrons.

Electrons are tiny, negatively charged particles that move around the nucleus. Unlike protons and neutrons, electrons have negligible mass but carry a negative charge (-1). They are arranged in specific energy levels called shells or orbitals, which surround the nucleus at increasing distances.

Each shell can hold a limited number of electrons: the first shell holds up to 2, the second up to 8, and the third up to 8 for the elements covered at this level. Electrons fill the shells starting from the one closest to the nucleus, moving outward as shells become full.

For example, a sodium atom (atomic number 11) has electrons arranged as 2 in the first shell, 8 in the second, and 1 in the third shell. This arrangement influences the atom's chemical properties and reactivity.

• Electrons have negative charge and negligible mass.
• Electrons occupy shells around the nucleus.
• Shells fill in order: 2, 8, 8 electrons for first three shells.
• Electron arrangement affects chemical behaviour.

The atom is electrically neutral because the number of protons (positive charges) equals the number of electrons (negative charges). Neutrons add mass and contribute to the stability of the nucleus by reducing repulsion between protons.

The arrangement of electrons in shells is crucial for chemical bonding and reactions. Atoms tend to be more stable when their outermost shell is full, which explains why atoms gain, lose, or share electrons during chemical reactions.

Understanding the structure of the atom helps explain the periodic table's organisation and the properties of elements.

• Atoms are neutral: protons = electrons.
• Neutrons stabilise the nucleus.
• Full outer electron shells confer stability.
• Electron arrangement explains chemical behaviour.

Helium has an atomic number of 2, meaning it has 2 protons in its nucleus. It also has 2 neutrons, giving a mass number of 4. Around the nucleus, helium has 2 electrons occupying the first shell, which can hold up to 2 electrons.

This full first shell makes helium chemically inert, as it does not need to gain or lose electrons to become stable. This example illustrates how the nucleus and electron shells define an atom's identity and chemical properties.

• Helium nucleus: 2 protons (+), 2 neutrons (neutral).
• 2 electrons fill the first shell completely.
• Full shell leads to chemical inertness.
• Demonstrates link between structure and properties.

Atoms consist of smaller particles called subatomic particles: protons, neutrons, and electrons. Each particle has distinct properties, including electric charge and mass, which determine the atom's overall structure and behaviour.

Protons and neutrons are located in the nucleus at the centre of the atom, while electrons orbit the nucleus in shells. Knowing their relative charges and masses helps explain atomic interactions and chemical reactions.

• Protons, neutrons, and electrons are the main components of atoms.
• Their charges and masses differ significantly.
• These differences influence atomic structure and properties.

Electric charge is a fundamental property that causes particles to attract or repel each other. Protons carry a positive charge, electrons carry a negative charge, and neutrons have no charge (neutral).

The relative charges are expressed in simple whole numbers for ease of comparison: protons have a charge of +1, electrons have a charge of -1, and neutrons have a charge of 0.

Because atoms have equal numbers of protons and electrons in their neutral state, their charges balance out, resulting in no overall charge.

• Proton charge: +1 (positive)
• Electron charge: -1 (negative)
• Neutron charge: 0 (neutral)

Mass is a measure of how much matter a particle contains. Protons and neutrons have similar masses, which are much greater than that of electrons.

The relative mass scale sets the proton's mass as 1 unit. Neutrons also have a relative mass of 1, while electrons have a much smaller relative mass of approximately 1/1836 (about 0.0005).

This means that nearly all the mass of an atom is concentrated in its nucleus, since electrons contribute very little to the total mass.

• Proton mass: 1 (relative unit)
• Neutron mass: 1 (relative unit)
• Electron mass: ~1/1836 (very small)

Because protons and neutrons have similar masses and are located in the nucleus, the nucleus contains almost all the atom's mass.

Electrons, being much lighter and negatively charged, orbit the nucleus in shells and determine the atom's size and chemical properties.

The balance of positive protons and negative electrons keeps the atom electrically neutral overall, unless electrons are gained or lost during chemical reactions.

• Nucleus contains protons and neutrons, accounting for most mass.
• Electrons orbit nucleus, determining size and chemical behaviour.
• Charge balance between protons and electrons maintains neutrality.

A hydrogen atom has one proton and one electron, with no neutrons. The proton has a relative charge of +1 and mass of 1, while the electron has a relative charge of -1 and a very small mass (~1/1836).

This simple structure illustrates how the positive charge of the proton is balanced by the negative charge of the electron, resulting in a neutral atom.

The mass of the hydrogen atom is almost entirely due to the proton, as the electron's mass is negligible.

• Hydrogen: 1 proton (+1 charge, mass 1)
• 1 electron (-1 charge, mass ~1/1836)
• No neutrons in most common hydrogen isotope

The proton number, also called the atomic number, is the total number of protons found in the nucleus of an atom. Since protons carry a positive charge, the proton number also equals the positive charge of the nucleus.

Each element has a unique proton number. For example, hydrogen has 1 proton, so its proton number is 1; carbon has 6 protons, so its proton number is 6. This uniqueness means the proton number defines the identity of the element.

• Proton number = number of protons in the nucleus.
• It is a whole positive integer.
• Defines the element's identity.

The proton number distinguishes one element from another. Atoms with the same proton number belong to the same element, regardless of the number of neutrons or electrons.

For example, all atoms with 8 protons are oxygen atoms. Even if the number of neutrons varies (forming isotopes), the element remains oxygen because the proton number is constant.

This is why the proton number is fundamental in chemistry: it tells us exactly which element we are dealing with.

• Same proton number = same element.
• Different proton numbers = different elements.
• Isotopes have the same proton number but different neutron numbers.

The periodic table arranges elements in order of increasing proton number from left to right and top to bottom.

This arrangement means the proton number determines an element's position in the periodic table. For example, sodium has proton number 11 and is placed after neon (proton number 10).

Knowing the proton number allows you to locate the element in the periodic table and predict its chemical behaviour based on its group and period.

• Elements ordered by increasing proton number in the periodic table.
• Proton number determines element's position.
• Position relates to chemical properties.

The proton number equals the number of electrons in a neutral atom, since atoms have no overall charge.

Electrons arrange themselves in shells around the nucleus, and the number of electrons depends on the proton number.

This electron arrangement determines how atoms interact chemically. For example, elements with the same proton number have the same electron configuration and similar chemical properties.

• Proton number = number of electrons in neutral atom.
• Electron arrangement depends on proton number.
• Chemical properties depend on electron arrangement.

An atom consists of a central nucleus surrounded by electrons in shells. The nucleus contains two types of subatomic particles: protons and neutrons. Protons carry a positive charge, while neutrons are neutral. Both have approximately the same mass, which is significantly greater than that of electrons. The number of protons defines the element, while neutrons contribute to the atom's mass but not its chemical identity.

• Nucleus contains protons (+) and neutrons (0).
• Electrons (-) orbit the nucleus in shells.
• Protons and neutrons have similar mass; electrons have negligible mass.

The mass number is the total number of protons and neutrons in an atom's nucleus. It is sometimes called the nucleon number because protons and neutrons are collectively known as nucleons. This number is always a whole number and is specific to each isotope of an element.

Mass number = number of protons + number of neutrons

• Mass number counts protons and neutrons only.
• Electrons are not included due to negligible mass.
• Mass number is a whole number specific to each isotope.

The atomic number is the number of protons in the nucleus and defines the element's identity. For example, all carbon atoms have 6 protons, so carbon's atomic number is 6. The mass number, however, varies between isotopes because the number of neutrons can differ.

Thus, atomic number = number of protons; mass number = protons + neutrons.

This distinction is crucial for understanding isotopes and atomic mass.

• Atomic number = number of protons; defines element.
• Mass number = protons + neutrons; varies between isotopes.
• Mass number is always equal to or greater than atomic number.

Isotopes are atoms of the same element with the same number of protons but different numbers of neutrons. This means isotopes have the same atomic number but different mass numbers.

For example, carbon-12 and carbon-14 are isotopes of carbon. Both have 6 protons, but carbon-12 has 6 neutrons (mass number 12), while carbon-14 has 8 neutrons (mass number 14).

Mass number helps to identify and distinguish isotopes, which can have different physical properties but similar chemical behaviour.

• Isotopes have same atomic number, different mass numbers.
• Mass number identifies specific isotopes.
• Example: Carbon-12 (6p + 6n), Carbon-14 (6p + 8n).

Consider an atom with 11 protons and 12 neutrons. To find its mass number, add the number of protons and neutrons:

Mass number = 11 + 12 = 23. This atom has a mass number of 23.

Since the atomic number is 11, this element is sodium (Na). The isotope is sodium-23.

• Mass number = protons + neutrons.
• Example: 11 protons + 12 neutrons = mass number 23.
• Isotope notation: Element symbol with mass number (e.g., Na-23).

Electronic configuration describes how electrons are arranged around the nucleus of an atom in different energy levels or shells. Each shell can hold a limited number of electrons, and these electrons occupy subshells (s, p, d, f) in a specific order. Understanding electronic configuration is essential because it influences an element's chemical behaviour and position in the Periodic Table.

For elements with atomic numbers 1 to 20, electrons fill the shells starting from the lowest energy level (closest to the nucleus) moving outward. The first shell (n=1) can hold up to 2 electrons, the second shell (n=2) up to 8, and the third shell (n=3) also up to 8 for these elements, with the fourth shell beginning to fill at element 19.

• Electronic configuration shows the distribution of electrons in shells and subshells.
• It determines chemical properties and bonding behaviour.
• Electrons fill shells in order of increasing energy.

The maximum number of electrons that can occupy a shell is given by the formula 2n², where n is the shell number (energy level). For example:

For n=1: 2 × 1² = 2 electrons maximum in the first shell.

For n=2: 2 × 2² = 8 electrons maximum in the second shell, and so on.

• Shell 1 holds up to 2 electrons.
• Shell 2 holds up to 8 electrons.
• Shell 3 holds up to 18 electrons in total, but for elements 1-20, only up to 8 electrons fill the third shell before the fourth shell begins to fill.

To determine the electronic configuration, start by placing electrons in the lowest energy shell and move outward until all electrons (equal to the atomic number) are placed.

For example: Element 2 (Helium) has 2 electrons, so its configuration is 2 (first shell full): 2

Element 8 (Oxygen) has 8 electrons: 2 in the first shell, 6 in the second shell, so configuration is 2,6 (or 1s² 2s² 2p⁴ in subshell notation). Element 3 (Lithium) has 3 electrons: 2 in the first shell, 1 in the second shell, configuration 2,1.

• Fill shells starting from the first shell.
• Count electrons equal to the atomic number.
• Example: Sodium (atomic number 11) configuration is 2,8,1.

Here are some examples for elements with atomic numbers 1 to 20:

Hydrogen (1): 1 electron → 1 configuration: 1 (1s¹)

Neon (10): 10 electrons → 2,8 (1s² 2s² 2p⁶) full first and second shells, making it chemically stable as a noble gas. - Calcium (20): 20 electrons → 2,8,8,2 (1s² 2s² 2p⁶ 3s² 3p⁶ 4s²) with electrons filling the fourth shell after the third shell is full.

• Hydrogen: 1 electron → 1
• Neon: 10 electrons → 2,8
• Calcium: 20 electrons → 2,8,8,2

The electronic configuration explains why elements in the same group have similar chemical properties: they have the same number of electrons in their outer shell. For example, all Group 1 elements have one electron in their outermost shell, making them highly reactive.

Knowing the electronic configuration also helps predict bonding types, ion formation, and placement in the Periodic Table.

For instance, elements with full outer shells (like noble gases) are chemically inert, while those with one or two electrons in their outer shell tend to lose electrons and form positive ions.

• Outer shell electrons determine chemical reactivity.
• Elements in the same group share outer electron numbers.
• Full outer shells correspond to chemical stability.

Noble gases are elements located in Group VIII (also known as Group 18) of the periodic table. These include helium (He), neon (Ne), argon (Ar), krypton (Kr), xenon (Xe), and radon (Rn).

A defining feature of noble gases is that their outermost electron shell is fully occupied by electrons. For example, neon has 8 electrons in its outer shell, which is the maximum for that shell in the second period. Helium is an exception with a full outer shell of 2 electrons, as its first shell can only hold 2.

This full outer shell configuration makes noble gases chemically stable and largely unreactive because they have little tendency to gain, lose, or share electrons to form chemical bonds. This stability is why noble gases are often called inert gases, although some heavier noble gases can form compounds under special conditions.

• Noble gases have full outer electron shells (e.g., neon: 2,8).
• Full outer shells confer chemical stability and low reactivity.
• Helium has a full outer shell with 2 electrons, as its first shell holds a maximum of 2.

Elements in Groups I to VII of the periodic table show a clear pattern: the group number corresponds to the number of electrons in their outermost shell. This pattern helps predict the chemical properties of elements.

For example, elements in Group I (alkali metals) have 1 electron in their outer shell, such as sodium (Na) with an electron configuration of 2,8,1. Group II elements have 2 outer electrons, like magnesium (Mg) with 2,8,2.

This relationship holds true for the main groups (also called representative elements) and is fundamental for understanding how elements react chemically, as the number of outer-shell electrons determines bonding behaviour.

• Group number (I to VII) = number of outer-shell electrons.
• Group I elements have 1 outer electron (e.g., sodium: 2,8,1).
• Group II elements have 2 outer electrons (e.g., magnesium: 2,8,2).
• This pattern helps predict element reactivity and bonding.

The period number of an element indicates how many electron shells are occupied in its atoms. Period 1 elements have electrons in only the first shell, period 2 elements have electrons in two shells, and so on.

For example, chlorine (Cl) is in period 3, so its electrons occupy three shells with the configuration 2,8,7. Calcium (Ca) is also in period 4, so it has four occupied shells with configuration 2,8,8,2.

This pattern reflects the filling of electron shells as atomic number increases and is important for understanding trends in atomic size and chemical properties across periods.

• Period number = number of occupied electron shells.
• Period 1 elements have electrons in 1 shell only (e.g., hydrogen: 1).
• Period 3 elements have electrons in 3 shells (e.g., chlorine: 2,8,7).
• Electron shells fill sequentially as period number increases.

Consider sodium (Na), which is in Group I and Period 3. It has 1 electron in its outer shell (group number 1) and electrons in 3 shells (period number 3). Its electron configuration is 2,8,1.

Chlorine (Cl) is in Group VII and Period 3. It has 7 electrons in its outer shell (group number 7) and 3 occupied shells (period number 3), with configuration 2,8,7.

Neon (Ne), a noble gas in Group VIII and Period 2, has a full outer shell of 8 electrons and 2 occupied shells (2,8). This full outer shell explains its chemical inertness.

• Sodium (Na): Group I → 1 outer electron; Period 3 → 3 shells; config 2,8,1.
• Chlorine (Cl): Group VII → 7 outer electrons; Period 3 → 3 shells; config 2,8,7.
• Neon (Ne): Group VIII → full outer shell (8 electrons); Period 2 → 2 shells; config 2,8.

Knowing the number of outer-shell electrons helps predict how an element will react chemically. Elements with 1 to 3 outer electrons tend to lose electrons and form positive ions, while those with 5 to 7 tend to gain electrons and form negative ions.

The full outer shell of noble gases means they rarely form ions or compounds, serving as a reference for chemical stability.

Understanding period numbers helps explain trends such as atomic radius decreasing across a period due to increasing nuclear charge pulling electrons closer, despite the same number of shells.

• Outer-shell electrons determine ion formation and bonding tendencies.
• Noble gases' full outer shells make them chemically inert.
• Period number relates to atomic size and periodic trends.