Chapter 1 - Elements and the Periodic Table

Elements and the Periodic Table

Atomic Number

• Definition: The number of protons in the nucleus of an atom.

• Each element has a unique atomic number.

• The periodic table is currently organized by increasing atomic number.

• Atoms have an equal number of protons and electrons, resulting in no overall charge.

• Ions are formed when atoms lose or gain electrons, acquiring a charge.

Atomic Structure

• Element: A pure substance containing only one type of atom.

• Atom: The smallest unit of matter.

• Nucleus: The central region of an atom, containing protons and neutrons.

• Proton: Positively charged particle within the nucleus.

• Neutron: Neutral particle within the nucleus.

• Electron: Negatively charged particle existing outside the nucleus.

• Subatomic particle: Particles found within the atom (protons, neutrons, electrons).

Key Terms

• Atomic number (n.): The number of protons in the nucleus of an atom.

• Chemical symbol (n.): An abbreviation used to represent a chemical element.

• Periodic table (n.): A table of chemical elements arranged by increasing atomic number.

• Molecule (n.): Two or more atoms bonded by sharing electrons (e.g., H2).

• Compound (n.): Two or more atoms of different elements bonded together (e.g., NH3).

• Ion (n.): An atom that has gained or lost electrons, resulting in a charge.

Mass Number

• Definition: Approximately equal to the sum of protons and neutrons in an atom.

• Electrons have negligible mass, so they are not considered in mass number calculations.

• Mass\ number = no.\ protons + no.\ neutrons

• The mass number is a relative value and does not have units.

• Calculating the number of neutrons:
  Number\ of\ neutrons = mass\ number - atomic\ number

Isotopes

• Definition: Variants of an element with the same atomic number but different numbers of neutrons.

• All atoms of a specific element have the same number of protons. For example, all carbon atoms have 6 protons.

• Isotopes of an element have different mass numbers (e.g., 12C, 13C, 14C).

• Isotopic notation (scientific notation) is used to represent isotopes.

Periodic Table (Part 1)

• The periodic table is an organizational tool to identify patterns and trends.

• Metals, non-metals, and metalloids.

• Electronic configurations (shell and subshell).

• Atomic radii.

• Relationships between structures and properties of elements. These properties includes: electronegativity, first ionisation energy, metallic and non-metallic character and reactivity.

Periods and Groups

• Periods: Rows in the periodic table.

• Groups: Columns in the periodic table.

• Elements are arranged by increasing atomic number (number of protons).

• Metals are generally on the left, non-metals on the right, and metalloids form a 'staircase' pattern.

• Elements in the same group have similar reactivity.

• Elements in the same period have the same number of electron shells.

• Group 1 (Alkali Metals) and Group 17 (Halogens) are highly reactive. Group 18 (Noble Gases) are generally inert.

Electron Shells

• Electrons are arranged in shells around the nucleus.

• The first shell (closest to the nucleus) holds a maximum of 2 electrons.

• Subsequent shells hold more electrons (8, 18, 32).

• Electron configuration: Arrangement of electrons in shells.

• Valence electrons: Electrons in the outermost shell.

• Valence shell: The outermost energy shell containing valence electrons.

• Energy shells: Orbits containing different levels of energy around the nucleus.

• Electron shell notation: Summary of electrons per shell (e.g., 2, 8, 1).

• Ground state: Electrons at their lowest possible energy level.

Blocks on the Periodic Table

• Elements are categorized into blocks based on the location of their valence electrons.

  • s-block: Groups 1 and 2.

  • p-block: Groups 13 to 18.

  • d-block: Groups 3 to 12.

  • f-block: Lanthanoids (atomic number 57–71) and actinoids (atomic number 89–103).

Electronic Subshell Configuration

• The Schrodinger model includes orbitals within each electron shell.

  • Orbitals: Regions with the highest probability of finding electrons.

• The Pauli Exclusion Principle states that each orbital can hold 0, 1, or 2 electrons.

• Electron subshell notation: Order of filling of sub-shells including s,p,d,f.

  • s subshell: 1 orbital, max. 2 electrons

  • p subshell: 3 orbitals, max. 6 electrons

  • d subshell: 5 orbitals, max. 10 electrons

  • f subshell: 7 orbitals, max. 14 electrons

• Aufbau principle: Rule that states that subshells are filled by electrons from the lowest to the highest energy level. The filling order is typically:
  1s < 2s < 2p < 3s < 3p < 4s < 3d < 4p < 5s < 4d < 5p < 6s < 4f < 5d < 6p < 7s ments like copper and chromium.

Core Charge

• Definition: The effective pull of the positive nucleus on the valence electrons.

• It increases across a period and is constant within a group.

• Core\ charge = number\ of\ protons - inner\ shell\ electrons

• Example: Lithium (Li) core charge = 3 - 2 = +1

• Example: Potassium (K) core charge = 19 - 18 = +1

Atomic Radii

• Definition: The distance from the center of an atom to the valence electrons.

• Atomic radius decreases across a period (left to right) and increases down a group.

• Core charge and the number of electron shells affect atomic radius.

• Across a period, increasing core charge pulls valence electrons closer, reducing the radius.

• Down a group, adding electron shells increases the atomic radius.

Periodic Table (Part 2)

• Trends in properties affected by periodicity: atomic radius, electronegativity, first ionization energy, metallic character, and reactivity.

• Periodicity: Characteristics of elements in a period.

Electronegativity

• Definition: The ability of an element to attract shared electrons in a chemical bond.

• Electronegativity increases across a period and up a group.

• Electronegativity increases with:

  • Increasing core charge

  • Decreasing atomic radius

  • Fewer occupied electron shells

• Fluorine is the most electronegative element; Noble gases aren’t considered.

First Ionization Energy

• Definition: The energy required to remove the first valence electron from an atom.

• First ionization energy increases across a period (due to increasing core charge and decreasing atomic radius).

• First ionization energy decreases down a group (due to increasing number of electron shells and greater distance of valence electrons from the nucleus).

Metallic Character

• The degree to which an element exhibits properties of a metal (luster, conductivity).

• Metallic character decreases across a period and increases down a group.

• Metals in groups 1 and 2 possess high metallic character (tendency to lose electrons).

• Non-metals possess high non-metallic character.

• Metalloids display a mixture of metallic and non-metallic properties.

Reactivity

• The tendency of an atom to lose or gain electrons.

• Noble gases (Group 18) are inert (unreactive).

• Among metals, Group 1 is the most reactive, followed by Group 2.

• Among nonmetals, Group 17 is the most reactive, followed by Group 16.

• Reactivity trends depend on how easily electrons are lost or gained, atomic radius, core charge, and electronegativity.

Recycling Critical Elements

• Critical elements: Elements in limited supply that could become depleted without recycling.

• Recycling is important for element recovery and reducing waste.

Critical Elements

• Help Protect PRAM:

  • Help - Helium

  • Protect - Phosphorus

  • P - Post-transition elements

  • R - Rare earth elements

  • A - And

  • M - Metalloids

• Examples: Helium, phosphorus, rare earth elements (REE), post-transition metals, metalloids.

Transitioning to a Circular Economy

• Sustainable development: Meeting present needs without compromising future generations.

• Linear economy: A 'take-make-dispose' model.

• Circular economy: A continuous cycle that focuses on the optimal use and re-use of resource from the extraction of raw materials through to the production of new materials.

• Recycling technological waste (e.g., electronics) reduces landfill and increases the supply of critical elements.

• This approach aligns with green chemistry principles, particularly waste prevention.

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