Ions, Atoms, and Isotopes

Ions, Atoms, and Isotopes

Learning Objectives

• Understand how to determine the number of protons, neutrons, and electrons in an ion.
• Learn to represent an atom's atomic number and mass number.
• Estimate the size and scale of atoms using SI units and the prefix 'nano'.
• Define isotopes.

What is an Ion?

• Atoms are neutral because they have an equal number of protons (positive charge) and electrons (negative charge).
• Atoms can lose or gain electrons.
  • If an atom gains one or more electrons, it becomes a negative ion (more electrons than protons).
  • If an atom loses one or more electrons, it becomes a positive ion (more protons than electrons).
• An ion is a charged atom or group of atoms.

Example: Oxygen Atom

• Oxygen has an atomic number of 8, so it has 8 protons ($8+$) and 8 electrons ($8-$-).
• If it gains two electrons, it has 8 protons ($8+$) and 10 electrons ($10-$-).
• The overall charge is 2-.
• Formula: $O^{2-}$
• The $O^{2-}$ ion has 8 protons, 8 neutrons, and 10 electrons.

Example: Lithium Atom

• If a lithium atom loses one electron, it has two electrons ($2-$-) and three protons ($3+$).
• It forms a positive ion with a single positive charge, $Li^+$.
• This $Li^+$ ion has 3 protons, 4 neutrons, and 2 electrons.

Representing Atomic Number and Mass Number

• Representation:
  • $^{mass \ number}_{atomic \ number}Symbol$
  • Example:
    • Carbon: $^{12}_6C$
    • Sodium: $^{23}_{11}Na$
• You can determine the number of protons, neutrons, and electrons:
  • Atomic number = number of protons = number of electrons.
  • Number of neutrons = mass number - atomic number.

Example: Sodium (Na)

• Atomic number = 11
• Mass number = 23
• Number of protons = 11
• Number of electrons = 11
• Number of neutrons = $23 - 11 = 12$

The Size of Atoms

• A person has about 7 billion, billion, billion atoms in their body, written as $7 \times 10^{27}$.
• Atoms are incredibly small and cannot be seen individually.
• An atom is about a tenth of a billionth of a meter across.

Atom Radius

• Radius: $1 \times 10^{-10} m$ or 0.1 nanometers (nm).

Nucleus Radius

• Radius: $1 \times 10^{-14} m$
• The nucleus occupies a very small space in the atom.
• Almost all of an atom is space, occupied by electrons.

Isotopes

• Atoms of the same element always have the same number of protons.
• Isotopes are atoms of the same element with different numbers of neutrons.
• Isotopes have the same atomic number but different mass numbers.
  • Example: Carbon-12 ($^{12}C$) and Carbon-13 ($^{13}C$).
    • Carbon-12 has 6 protons and 6 neutrons.
    • Carbon-13 has 6 protons and 7 neutrons.
• Extra neutrons can make the nucleus unstable, leading to radioactivity.
• Not all isotopes are radioactive.
• Different isotopes have different physical properties (e.g., density) but the same chemical properties.
  • Chemical properties depend on electronic structures.
  • Isotopes have the same number of protons and electrons, so the electronic structure is the same.

Hydrogen Isotopes

• Hydrogen (Hydrogen-1): $^1H$
• Deuterium (Hydrogen-2): $^2H$
• Tritium (Hydrogen-3): $^3H$
• Each has a different mass; tritium is radioactive, but all have identical chemical properties.
  • $2H<em>2(g) + O(g) \rightarrow 2H</em>2O(l)$

Key Points

• Atoms that gain electrons form negative ions; atoms that lose electrons form positive ions.
• Atomic number and mass number representation: $^{24}_{12}Mg$
• Isotopes are atoms of the same element with the same number of protons but different numbers of neutrons; they have identical chemical properties but different physical properties.

Electronic Structures

Learning Objectives

• Understand how electrons are arranged in an atom.
• Know the electronic structures of the first 20 elements in the Periodic Table.
• Represent electronic structures in diagrams and using numbers.

Electron Arrangement

• Electrons are arranged around the nucleus in shells, representing different energy levels.
• The lowest energy level is the shell nearest to the nucleus.
• Electrons occupy the lowest available energy level.

Electron Shell Diagrams

• Each energy level (shell) can hold a specific number of electrons:
  • The first energy level can hold up to 2 electrons.
  • The second energy level can hold up to 8 electrons.
  • The third energy level can hold up to 8 electrons before the fourth starts to fill.
• Beyond the first 20 elements, the situation becomes more complex.

Example: Sodium Atom

• Atomic number of 11, so it has 11 protons and 11 electrons.
• Electronic structure: 2,8,1.

Electronic Structure Notation

• Write down the numbers of electrons in each energy level, separated by commas.

Example: Silicon

• 14 electrons, in Group 4 of the Periodic Table.
• Electronic structure: 2,8,4.

Electronic Structure and the Periodic Table

• Elements in the same group have the same number of electrons in their highest energy level (outer electrons).

Example: Group 1 Elements

• Lithium, sodium, and potassium each have one electron in their outermost shell.

• This makes them very reactive.

• The chemical properties of an element depend on the number of electrons in its highest energy level (outermost shell).

• Elements in the same group react similarly because they have the same number of electrons in their highest energy level.

Group 1 Elements Reactivity with Water

• Lithium + water → lithium hydroxide + hydrogen
• Sodium + water → sodium hydroxide + hydrogen
• Potassium + water → potassium hydroxide + hydrogen

Noble Gases (Group 0)

• Very unreactive elements.
• Atoms have a stable arrangement of electrons with eight electrons in the outer shell, except for helium, which has two.

Key Points

• Electrons in an atom are arranged in energy levels or shells.
• The first shell can hold up to 2 electrons, and the second shell can hold up to 8 electrons. The 4th shell starts after 8 electrons occupy the 3rd shell.
• The number of electrons in the outermost shell determines how an element reacts.