Chemistry Unit 8: Importance of Electrons

Unit 8 – Why do electrons matter?

• Unit 8 explores the significance of electrons in various chemical contexts.

Unit 7 Recap

• Atoms consist of protons, neutrons, and electrons.
• Atomic mass (AR) is calculated as the sum of protons and neutrons.
• Alpha decay is a type of radioactive decay where an atomic nucleus emits an alpha particle, causing a change in the mass number and atomic number of the decaying nucleus.

Unit Outline

• Factual: The main energy level influences atomic reactivity and electron arrangement representation is important.
• Conceptual: Electron shell organization determines an atom's oxidation state.
• Debatable: The ability to control reduction-oxidation of metals could define human progress.

Atomic Structure

• Physicists made significant advancements in understanding atomic structure.
• Thomson proposed the plum pudding model.
• Rutherford suggested atoms have a dense nucleus.
• Bohr described electrons as occupying orbits around the nucleus.
• Electrons can absorb specific energy levels (quantum).
• The Schrodinger wave equation is the most accurate model to date.

Atomic Structure (continued)

• Electrons can be excited and release energy when decaying.
• Different elements absorb different amounts of energy due to distinct energy level variations.

Energy level and atom's reactivity

• Create two tables in your notebook.
• The first table should contain a compound name, chemical formula and observations.
• The second table should detail wavelength, colour interval/nm, and frequency interval/THz associated with different colours.

Flame Test – Discussing Observations

• Investigate the reason behind colour changes observed in metal and non-metal flame tests.
• Analyze whether the effect is due to a physical change or chemical reaction.
• Use the table to associate particular wavelengths with colours.
• Evaluate the applicability of this method/test in fireworks.

Emission Spectra

• Wavelength and frequency are inversely proportional.
• The speed of light (c) is 3.00 x 10⁸ m/s.
• Frequency is denoted by (f) and is also related to energy (E=hxf).
• Planck's constant (h) is 6.63 x 10⁻³⁴ Js.

Unit 7 Review - Practice Problems

• Complete problems on page 164 numbered 1, 3, 4, 6, 7, 9, 11, and 14.

Atomic radius

• Atoms increase in size with increasing electron addition.
• Different periodic table elements and groups have varied atomic radii sizes.

Octet rule

• Electrons prioritize orbits closest to the nucleus.
• Circular disks represent energy levels.
• The first energy level holds 2 electrons and the second holds 8.

Electron arrangement (EA)

• Electron arrangement (EA) facilitates electron configuration representation.
• Group electrons together, separating them with commas.
• Use resources, including videos like https://www.youtube.com/watch?v=hSkJzE2Vz_w, to assist in understanding electron arrangements.
• Elements in the same column/group share the same number of valence (outermost) electrons.

Patterns

• The outermost shell is called the valence shell.
• It contains valence electrons.
• Noble gas elements (excluding He) generally have 8 valence electrons.
• 2n² describes the maximum number of electrons in any shell where n is the shell number.
• Shells aren't always filled immediately to their full capacity.
• Inner shell electrons are referred to as core electrons.

Atomic radius (continued)

• Atomic radius decreases across periods.
• This relates to increasing nuclear charge.
• Increased shielding from core electrons plays a role as well.
• Elements show increased radii down groups.

Ionization Energy

• Ionization energy (IE) is the energy needed to remove an electron from a neutral atom.
• Ionization energy decreases down the groups and increases across periods.
• Outermost electrons are held by the nucleus and their proximity influences the energy required for removal.

Electronegativity (x)

• Electronegativity helps predict bond types (ionic or covalent).
• Electronegativity is the ability of an atom to attract electrons in a chemical bond.
• Differences in electronegativity values predict the nature of the chemical bond between elements. Higher difference in electronegativity value means ionic bonding.

Pauling scale

• Electronegativity values are dimensionless and range from 0.8 to 4.0.
• Fluorine has the highest electronegativity in the Pauling scale.
• Caesium has the lowest electronegativity in the Pauling scale.
• The larger the difference in electronegativity between atoms, the more ionic the bond character.
• Ionic bonding occurs when the electronegativity difference between atoms is greater than 1.8.

Valence

• Elements within the same group share a similar number of valence electrons.
• Some elements lose electrons and others gain electrons to form ions.
• The oxidation number of an atom in a chemical compound represents the number of electrons lost or gained by an atom during a chemical reaction.

Valency (continued)

• Valency is related to the oxidation number (but not the same thing).
• Transition metals have variable valency.
• Copper forms copper(I) and copper(II) ions.
• Electron arrangement notation only works well for a particular number of elements.
• The third electron shell can hold up to 18 electrons.

Dot and cross diagrams

• Diagrams to represent electron shell expansion.
• Used to show how atoms form bonds and compounds.
• Describe how atoms of differing valencies form compounds.
• Use OILRIG: Oxidation Is Loss, Reduction is Gain
• Show electrons lost or gained when forming compounds.

Dot-and-cross diagrams for ionic compounds

• Brackets denote ionic bonds.
• Electrons are not shared in ionic compounds.

Diagrams for multiple covalent bonds

• Double and triple bonds are represented by four and six closely grouped electrons.

Lewis structures

• Alternative method to depict the octet rule.
• Focus on valence electrons.
• All elements within the same group have the same number of valence electrons.
• Lewis structures use dots to represent valence electrons.

Lewis structures (Continued)

• Paired electrons are represented as two dots.
• Non-bonding electrons are referred to as "lone pairs."
• Unpaired electrons are shown as single dots.
• Brackets are used for ionic bonds in Lewis structure representations.

Water - Worked Example

• A detailed step by step guide/worked example to draw a Lewis structure for water (H₂O).

Redox reactions with Lewis structures

• Redox reactions involve the transfer of electrons.
• Oxidation involves the loss of electrons, reduction involves the gain of electrons.
• OILRIG is a helpful acronym for remembering the terminology.

Homework - Practice Lewis structures

• Draw Lewis structures for water (H₂O), Nitrogen trichloride (NCI₃) and Carbon dioxide (CO₂).

Practicing Lewis Structures

• Instructions and information to practice drawing Lewis structures are given.