Atomic structure and the Periodic Table Revision Notes

The nucleus is the dense central core of an atom. It contains two types of subatomic particles: protons and neutrons.

Protons carry a positive electrical charge (+1) and have a relative mass of approximately 1 atomic mass unit (amu).

Neutrons have no electrical charge (neutral) and a similar mass to protons, about 1 amu. Together, protons and neutrons are called nucleons because they reside in the nucleus.

• Nucleus contains protons and neutrons.
• Protons are positively charged; neutrons are neutral.
• Nucleus accounts for nearly all the atom's mass.

Electrons are negatively charged particles that orbit the nucleus in defined energy levels called shells.

Each shell can hold a limited number of electrons: the first shell holds up to 2 electrons, the second up to 8, and the third up to 8 (for elements covered at this level).

Electrons have negligible mass compared to protons and neutrons but determine the atom's chemical properties and bonding behaviour.

• Electrons orbit nucleus in shells.
• Shells have maximum electron capacities (2, 8, 8).
• Electrons are involved in chemical reactions.

An atom is electrically neutral because it contains equal numbers of protons and electrons.

The nucleus is very small compared to the overall size of the atom but contains almost all the mass.

Electrons move rapidly in shells around the nucleus, creating a cloud of negative charge that balances the positive charge of the nucleus.

• Atoms have equal protons and electrons for neutrality.
• Nucleus is tiny but massive.
• Electron shells define atomic size.

The number of protons (atomic number) defines the element and its position in the Periodic Table.

Neutrons contribute to the mass number and can vary in isotopes of the same element.

Electron arrangement in shells influences chemical reactivity and bonding patterns.

• Proton number identifies the element.
• Neutrons affect atomic mass and isotopes.
• Electron shells determine chemical behaviour.

Atoms consist of three primary 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. The relative charges and masses of these particles influence atomic interactions and chemical reactions.

• Protons: positively charged particles in the nucleus.
• Neutrons: neutral particles in the nucleus.
• Electrons: negatively charged particles orbiting the nucleus.

The electric charge of subatomic particles is measured relative to the elementary charge, which is the charge of a proton or electron.

Protons carry a positive charge, neutrons have no charge, and electrons carry a negative charge equal in magnitude to that of protons but opposite in sign.

This can be summarised as: +1 for protons, 0 for neutrons, and -1 for electrons (relative charge units).

• Proton charge = +1 (relative)
• Neutron charge = 0 (neutral)
• Electron charge = -1 (relative)

The mass of subatomic particles is often expressed relative to the mass of a proton, which is assigned a relative mass of 1.

Protons and neutrons have approximately the same mass, each with a relative mass of 1.

Electrons are much lighter, with a relative mass of about 1/1836, which is often approximated as zero in many calculations due to its negligible contribution to atomic mass.

• Proton mass = 1 (relative)
• Neutron mass = 1 (relative)
• Electron mass ≈ 1/1836 (relative)

The positive charge of protons balances the negative charge of electrons, making atoms electrically neutral overall when the numbers of protons and electrons are equal.

The mass of an atom is concentrated in the nucleus because protons and neutrons have much greater mass than electrons.

The relative masses and charges determine the chemical properties of elements and how atoms interact to form compounds.

• Atoms are neutral if proton number equals electron number.
• Nucleus contains nearly all the atom's mass.
• Electron charge influences chemical bonding and reactions.

Calculate the relative atomic mass of an atom with 6 protons, 6 neutrons, and 6 electrons.

Since electrons have negligible mass, only protons and neutrons contribute significantly to atomic mass.

Relative atomic mass = number of protons × relative mass of proton + number of neutrons × relative mass of neutron + number of electrons × relative mass of electron = (6 × 1) + (6 × 1) + (6 × 1/1836) ≈ 6 + 6 + 0.0033 ≈ 12.0033, which is approximately 12.

• Relative atomic mass ≈ 12
• Electrons contribute negligible mass

The proton number, also known as the atomic number, is defined as the total number of protons present in the nucleus of an atom. Protons are positively charged particles located in the nucleus, and their number is unique to each chemical element.

For example, hydrogen has a proton number of 1 because it has one proton in its nucleus, while carbon has a proton number of 6, indicating six protons in its nucleus.

• Proton number = number of protons in the nucleus.
• Also called atomic number.
• Unique for each element.

The proton number determines the identity of an element. No two different elements have the same proton number. This means that the proton number is the defining characteristic of an element.

Because the proton number is unique, it is used to arrange elements in the Periodic Table in order of increasing atomic number. This arrangement reflects periodic trends in chemical properties.

For example, all atoms with proton number 8 are oxygen atoms, regardless of their isotope or ionisation state.

• Proton number uniquely identifies an element.
• Elements are arranged in the Periodic Table by increasing proton number.
• Chemical properties depend on proton number.

In a neutral atom, the number of electrons equals the number of protons to balance the positive and negative charges, resulting in no overall charge.

Electrons orbit the nucleus in shells, and their arrangement is influenced by the proton number because the positive charge of the nucleus attracts electrons.

For example, a neutral sodium atom has 11 protons and 11 electrons.

• Neutral atom: number of electrons = number of protons.
• Electron arrangement depends on proton number.
• Charge neutrality is maintained by equal protons and electrons.

The proton number influences the chemical behaviour of an element because it determines the number of electrons and their arrangement in shells around the nucleus.

Chemical reactions involve the electrons, especially those in the outermost shell (valence electrons). Since the proton number fixes the electron configuration, it indirectly controls reactivity and bonding.

For example, elements in the same group of the Periodic Table have the same number of outer shell electrons and similar chemical properties because they have the same proton number pattern.

• Proton number determines electron configuration.
• Electron configuration controls chemical reactivity.
• Elements with similar proton number patterns have similar properties.

The mass number, often referred to as the nucleon number, is defined as the total count of protons and neutrons present in the nucleus of an atom. Protons and neutrons are collectively called nucleons because they reside in the nucleus.

Mass number is a whole number and is symbolised by the letter A in nuclear notation. It is distinct from the atomic number (Z), which counts only the protons in the nucleus.

• Mass number (A) = number of protons + number of neutrons
• Protons and neutrons are nucleons located in the nucleus
• Mass number is always a whole number

The mass number provides information about the total mass of the nucleus since protons and neutrons have approximately equal masses, each close to 1 atomic mass unit (amu). Electrons have negligible mass compared to nucleons and do not contribute significantly to the mass number.

Knowing the mass number helps chemists distinguish between different isotopes of the same element. Isotopes have the same number of protons (same atomic number) but different numbers of neutrons, resulting in different mass numbers.

• Mass number approximates the atomic mass of the nucleus
• Electrons have negligible mass and are excluded from mass number
• Isotopes differ in neutron number and thus in mass number

Atoms are often represented using nuclear notation, which shows the mass number (A) as a superscript and the atomic number (Z) as a subscript to the left of the chemical symbol.

For example, the isotope carbon-12 is written as C, where 12 is the mass number (6 protons + 6 neutrons) and 6 is the atomic number (number of protons).

This notation clearly indicates the composition of the nucleus in terms of protons and neutrons.

• Nuclear notation: C (mass number superscript, atomic number subscript)
• Carbon-12 has 6 protons and 6 neutrons
• Mass number = 6 + 6 = 12

Calculate the mass number of an atom with 17 protons and 20 neutrons.

Step 1: Identify the number of protons = 17 (atomic number)

Step 2: Identify the number of neutrons = 20 (given) Step 3: Calculate mass number = protons + neutrons = 17 + 20 = 37 Step 4: Write the nuclear notation: 7
17Cl (Chlorine-37 isotope)

• Mass number = protons + neutrons
• Example: 17 + 20 = 37
• Nuclear notation: 7
  17Cl

The atomic number (Z) is the number of protons in the nucleus and defines the element's identity. The mass number (A) includes both protons and neutrons and varies among isotopes of the same element.

For example, chlorine has two common isotopes: chlorine-35 and chlorine-37. Both have 17 protons (atomic number 17) but differ in neutron number (18 and 20 respectively), resulting in different mass numbers.

It is important not to confuse these two numbers when describing atoms or writing nuclear notation.

• Atomic number = number of protons (defines element)
• Mass number = protons + neutrons (defines isotope)
• Isotopes have same atomic number but different mass numbers

Electronic configuration describes how electrons are arranged in an atom's electron shells or energy levels. 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 elements with atomic numbers up to 20. The number of electrons in a neutral atom equals its proton number (atomic number). Electrons fill shells starting from the lowest energy level (closest to the nucleus) following the Aufbau principle.

The notation for electronic configuration lists the number of electrons in each shell separated by commas. For example, 2,8,1 means 2 electrons in the first shell, 8 in the second, and 1 in the third. This notation is standard for elements with atomic numbers up to 20.

• First shell maximum: 2 electrons
• Second shell maximum: 8 electrons
• Third shell maximum (up to element 20): 8 electrons
• Number of electrons = proton number for neutral atoms

Three main principles guide electron arrangement: the Aufbau principle states electrons occupy the lowest energy orbitals first; the Pauli exclusion principle states each orbital holds a maximum of two electrons with opposite spins; Hund's rule states electrons fill degenerate orbitals singly before pairing. For elements with proton numbers 1 to 20, these rules result in filling shells in order: 1s, 2s, 2p, 3s, and 3p orbitals.

For example, oxygen (atomic number 8) has the configuration 2,6, meaning 2 electrons in the first shell and 6 in the second shell, filling the 2s and 2p orbitals accordingly.

• Aufbau principle: fill lowest energy orbitals first
• Pauli exclusion principle: max two electrons per orbital with opposite spins
• Hund's rule: fill orbitals singly before pairing

To calculate the electronic configuration of a neutral atom:

1. Identify the atomic number (number of protons and electrons).

2. Fill the first shell with up to 2 electrons, then the second shell with up to 8, then the third shell with up to 8 (for elements up to atomic number 20).

• Example: Sodium (Na), atomic number 11
• First shell: 2 electrons
• Second shell: 8 electrons
• Third shell: 1 electron
• Electronic configuration: 2,8,1

Ions form when atoms gain or lose electrons to achieve a full outer shell, often resembling the nearest noble gas configuration. To find the electronic configuration of an ion:

1. Determine the charge of the ion (positive means electrons lost, negative means electrons gained).

2. Adjust the number of electrons accordingly from the neutral atom configuration, removing or adding electrons from/to the outermost shell first.

• Example: Sodium ion (Na⁺)
• Neutral Na: 2,8,1
• Na⁺ loses 1 electron: 2,8
• Example: Oxide ion (O²⁻)
• Neutral O: 2,6
• O²⁻ gains 2 electrons: 2,8

Magnesium (Mg) has atomic number 12. Calculate its electronic configuration and that of Mg²⁺ ion.

Step 1: Neutral Mg has 12 electrons. Fill shells: first shell 2, second shell 8, third shell 2. Configuration: 2,8,2.

Step 2: Mg²⁺ ion has lost 2 electrons, so it has 10 electrons. Remove electrons from outer shell: 2,8. Configuration: 2,8.

• Mg: 2,8,2 (neutral)
• Mg²⁺: 2,8 (ion)

Forgetting to adjust electron count when calculating ions.

Adding or removing electrons from inner shells instead of the outer shell.

Miscounting electrons in shells beyond the maximum capacity (e.g., more than 8 in second shell).

• Always start electron removal or addition from the outermost shell.
• Check total electrons equal atomic number minus or plus ion charge.
• Remember shell capacities: 2 for first, 8 for second and third (up to element 20).

The periodic table arranges elements in order of increasing proton number (atomic number). Elements are organised into periods (rows) and groups (columns). Each period corresponds to the number of electron shells occupied by electrons in the atoms of the elements in that period.

For example, elements in Period 1 have electrons in only one shell, Period 2 elements have electrons in two shells, and so on. This pattern reflects the filling of electron shells as atomic number increases.

Groups are vertical columns where elements share similar chemical properties because they have the same number of electrons in their outermost shell (valence electrons). This similarity in outer shell electron number influences their chemical reactivity and bonding behaviour.

• Period number = number of occupied electron shells.
• Group number (for Groups I to VII) = number of electrons in the outer shell.

Noble gases are elements in Group VIII (also called Group 0 or 18 in some tables). They have a full outer electron shell, which makes them very stable and chemically unreactive under normal conditions.

For example, helium has 2 electrons filling its first shell, neon has 8 electrons filling its second shell, and argon has 8 electrons filling its third shell. This full outer shell configuration means noble gases have little tendency to gain, lose or share electrons, explaining their inertness.

The general electron configuration for noble gases ends with a complete outer shell: 2 electrons for Period 1 (helium), and 8 electrons for Periods 2 and beyond (neon, argon, etc.).

• Noble gases have complete outer shells: 2 electrons for helium, 8 for others.
• Complete outer shell confers chemical stability and low reactivity.

For Groups I to VII, the group number corresponds to the number of electrons in the outermost shell of the element's atoms.

This means: Group I elements have 1 outer electron, Group II have 2, Group III have 3, and so forth up to Group VII which have 7 outer electrons.

This pattern explains the chemical properties and valency of elements in these groups. For example, Group I metals (alkali metals) have 1 electron to lose to form +1 ions, while Group VII elements (halogens) have 7 outer electrons and tend to gain 1 electron to form -1 ions.

• Group I elements: 1 outer electron.
• Group II elements: 2 outer electrons.
• Group III elements: 3 outer electrons.
• Group IV elements: 4 outer electrons.
• Group V elements: 5 outer electrons.
• Group VI elements: 6 outer electrons.
• Group VII elements: 7 outer electrons.

The period number of an element indicates how many electron shells are occupied by electrons in its atoms.

For example, sodium (Na) is in Period 3, so it has electrons in three shells: 2 in the first shell, 8 in the second, and 1 in the third.

This relationship helps predict the electron configuration of elements and understand their chemical behaviour.

• Period 1 elements have electrons in 1 shell.
• Period 2 elements have electrons in 2 shells.
• Period 3 elements have electrons in 3 shells, etc.

Sulfur (S) is in Group VI and Period 3 of the periodic table.

Step 1: Determine the number of occupied shells from the period number: Period 3 means 3 shells.

Step 2: Determine the number of outer shell electrons from the group number: Group VI means 6 outer electrons in the outermost shell (third shell). The total number of electrons in sulfur is 16 (atomic number). The first shell holds 2 electrons, the second shell holds 8 electrons, and the remaining 6 electrons are in the third shell (outer shell).

• Electron configuration of sulfur: 2, 8, 6.
• Number of occupied shells = 3 (Period 3).
• Number of outer electrons = 6 (Group VI).