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.
    • 2H2(g) + O(g) \rightarrow 2H2O(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.