Atoms, Elements and Compounds

Unit 1: Diversity of Materials

• Unit 1 explores how to explain the diversity of materials.
• It addresses how the chemical structures of materials explain their properties and reactions.
• It looks at how materials are quantified and classified.

1A Atoms and Elements

• Elements, isotopes, and ions are defined, including appropriate notation.
• Notation includes atomic number, mass number, and the number of protons, neutrons, and electrons.
• All matter consists of atoms which are tiny particles with a diameter of 0.1 - 0.5 nanometres
• 1 nanometre (nm) is one-billionth of a meter, or 1 x 10-9 m

Chemical Elements and Compounds

• A chemical element is a substance made of only one type of atom.
• Atom arrangement varies by type of element, examples:
  • Copper (Cu) is a metallic lattice.
  • Carbon (C) is covalent network lattice.
  • Sulfur (S8) is a molecule.
• A compound is made of different types of atoms.
• Compounds consist of non-metal atoms (molecules) or non-metal and metal atoms (ionic lattices).
• Compounds can be broken into simpler substances (atoms), but elements cannot.

Atomic Components

• Atoms consist of:
  • Nucleus: Contains protons and neutrons.
  • Protons: Positively charged particles (+1).
  • Neutrons: Neutral particles (0).
  • Electrons: Negatively charged particles (-1) that orbit the nucleus.
• Relative size and mass:
  • Protons have a relative size of 1 and a mass of 1.6735 × 10^-27 kg.
  • Neutrons have a relative size of 1 and a mass of 1.675 x 10^-27 kg.
  • Electrons have a relative size of 1/1800 and a mass of 9.109 x 10^-31 kg.

Atomic and Mass Numbers

• Atomic number defines the number of protons in an atom of an element.
• Mass number indicates the number of protons and neutrons in an element.
• The mass number is a relative value with no units.

Isotopes

• The number of protons defines the element.
• Isotopes have the same number of protons but different numbers of neutrons.

1B The Periodic Table (1)

• The periodic table is an organizational tool used to identify patterns, trends, and relationships between elements.
• This includes structures (shell and subshell electronic configurations and atomic radii) and properties.
• Properties include electronegativity, first ionization energy, metallic/non-metallic character, and reactivity.

Electron Configurations

• Electron configurations involve shells and subshells.
• Schrodinger proposed a refinement on the Bohr model where electrons don't occupy orbits but regions called orbitals
• Pauli exclusion principle applies, which states that there can be only 2 electrons per orbital
• Shells or energy levels also have "subshells."
• Electrons fill the lowest energy levels first, also known as the Aufbau principle.

Shells, Subshells, and Orbitals

• The maximum number of electrons in a shell is:

  • Shell 1: 2 electrons
  • Shell 2: 8 electrons
  • Shell 3: 18 electrons
  • Shell 4: 32 electrons

• Shells and subshells around the nucleus are by increasing energy:

  • 1s , 2s , 2p, 3s, 3p, 4s , 3d, 4p.

• The different shells, subshells and orbitals in an atom include:

  • Number of shell 1: 1s subshell(s) in shell, 1 Number of orbitals in subshell and a Maximum number of electrons of 2
  • Number of shell 2: 2s and 2p subshell(s) in shell, 1 and 3 Number of orbitals in subshell and a Maximum number of electrons of 2 and 6 respectivley
  • Number of shell 3: 3s, 3p and 3d subshell(s) in shell, 1, 3 and 5 Number of orbitals in subshell and a Maximum number of electrons of 2, 6, and10 respectivley
  • Number of shell 4: 4s, 4p, 4d and 4f subshell(s) in shell, 1, 3, 5 and 7 Number of orbitals in subshell and a Maximum number of electrons of 2, 6, 10 and 14 respectivley

• Example electron configurations include: 1s22s22p63s23p4

Exceptions to Electron Configuration Rules

• For Chromium, 3d5, 4s¹ is slightly more stable than fully filling the 4s shell with 2 electrons before adding to 3d.
• For Copper, 3d10, 4s¹ is slightly more stable than fully filling the 4s shell with 2 electrons before adding 9 electrons.

Ground State vs. Excited State

• Ground State: All electrons are in the lowest possible shells/subshells.
• Excited State: One or more electrons have been promoted to a higher shell/subshell by absorbing energy.

Atoms vs. Ions

• Electronic configuration changes when atoms become ions due to gaining or losing electrons.

Periodic Table Organization

Rows (periods) indicate the same number of occupied shells.

• Columns (groups) indicate the same number of valence electrons.

Trends in Atomic Radius

• The size of the atom largely depends on the valence shell
• The number of protons increases as we move across a period.
• Atomic radius decreases across a period.

Core Charge

• Described as the atom as a positive nucleus, inner shell electrons and valence shell electrons
• The atom as a whole is neutral, inner shell electrons shield valence shell electrons from the nucleus therefore affect the size of the atom.
• How the core charge varied along the periodic table as you go across a period is that they decrease.

Trends Down a Group

• As you go down a group the arrangement explains the increase in atomic radius because they number of shells increases

Periodic Table Blocks

• The periodic table is divided into blocks that reflect the highest energy subshell found in the atoms of an element.
• The four blocks are s, p, d, and f.
• Groups (columns) of elements have the same number of outer valence electrons.
• Periods (rows) of elements have the same number of occupied energy shells.
• Atomic radius decreases across a period and increases down a group.

1C The Periodic Table (2)

• Focuses on trends in electronegativity, ionization energy, metallic character, and reactivity.
• These properties depend on core charge and atomic radius.

Electronegativity

• Electronegativity measures how strongly an atom attracts electrons towards itself.
• It is measured on the Pauling Scale.
• Highly electronegative elements strongly attract electrons.
• Electronegativity varies across a period due to the change in the size of atomic radii.
• Electronegativity decreases down a group.

Ionization Energy

• Ionization energy refers to why some elements lose their electrons more easily than others.
• Greater core charge results in a stronger attraction for electrons and makes it harder to ionize.
• A larger atom results in less attraction and easier ionization.

Trends Affecting Electronegativity and Ionization Energy

• How the core charge effects Electronegativity
• How the atomic radius effects Electronegativity
• First lonisation Energy describes the energy that causes the atom to lose electrons

Metallic and non Metallic character

• Metals readily lose electrons, which gives rise to many of their properties.
• This is explained by atomic radius and core charge.
• Metalloids are conductive but not malleable and ductile.
• Metallic character decreases across a period and increases down a group.
• Non-metals accept electrons to achieve a full valence shell, making them useful in creating molecules

Reactivity Trends

• Reactivity is how easy it is for an atom to gain or lose electrons.
• Reactive metals donate electrons. non-metals accept electrons.
• A large atomic radius increases the ability to lose electrons but reduces the ability to attract electrons due to shielding electrons.
• An increase in core charge allows atoms to better hold onto their valence electrons and gain electrons more easily.

1D Critical Elements

• Critical elements are those that are both endangered and vital to humans.

Examples of Critical Elements and Their Uses

• Helium: Used in party balloons and MRI scanners.
• Phosphorus: Used in fertilizers and in the form of phosphate salts.
• Rare-earth: Have many uses and are mostly from the lanthanoids on the periodic table, difficult to extract from rocks.
• Transition metals: Are used heavily in electronics with Examples include iridium, platinum, osmium and palladium.
• Indium: Is a metalloid used in touch screens and automotive glass, also used in solar panels

Ways to make elements more sustainable

• Increase the amount or reduce the need for the elements
• Improve methods of recycling the elements