Atomic Structure and Notation

Atomic Structure

• Electrons possess a negative charge of -1.
• Electrons reside in the electron cloud (orbitals) and have a relative mass of 1/1836.
• The electron cloud (orbitals) is a region of space where electrons are likely to be found, given their small size and wave-like properties which prevent determination of both exact speed and location.
• Atoms are the smallest unit of an element retaining its composition and consist mostly of empty space occupied by light, negatively charged electrons.
• Protons have a positive charge of +1.
• Protons define the element of an atom.
• Protons are tightly packed into the nucleus and only vibrate at a fixed position, possessing a relative mass of one.
• Neutrons have a neutral charge.
• Neutrons are tightly packed into the nucleus and only vibrate about their fixed position, possessing a relative mass of one.

Notation

• Nuclear notation displays the mass number (A) and atomic number (Z) along with the element symbol (X).
• A represents the mass number, which equals the number of protons plus neutrons.
• Z signifies the atomic number, representing the number of protons in the nucleus.
• Atoms are neutrally charged, indicating the number of protons equals the number of electrons.
• Ions have a charge (positive or negative) where the number of protons is not equal to the number of electrons.
• Cations are ions with more protons than electrons.
• Anions are ions with fewer protons than electrons.
• Ions can achieve ‘Noble Gas Configurations’.
• Isotopes share the same proton number but vary in neutron numbers.
• Isotopes possess similar chemical properties but different physical properties, such as mass.

Representing Electron Arrangements

• Bohr diagrams show all electrons and electron shells within an atom.
• Lewis diagrams represent valence (outermost) electrons only.
• Electron configurations list the number of electrons in each shell, moving from inner to outer shells (e.g., carbon: 2,4).

Electron Arrangement

• Electrons exist in layers called shells or energy levels.
• Electron configuration describes the arrangement of electrons in these shells.
• Atoms are mostly empty space.
• Each shell has a maximum electron capacity, with electrons filling shells closest to the nucleus first.

Energy Levels

• Electrons occupy different shells, each having different energy levels.
• Varying energy levels can be observed through flame isolation experiments and emission spectra as different colours.
• An electron's energy level depends on its energy amount.
• Shells increase in energy from the innermost shell outward.
• Valence electrons can become 'excited' (move to higher energy levels) when heated by a flame.
• When an electron returns to its original energy level, or "ground state," it emits visible light.
• Ground state electrons absorb energy to move to a higher energy level.
• An unstable excited state electron returns to the ground state, emitting a specific wavelength of light energy.

Valence Electrons

• Metals that lose valence electrons more easily are more reactive.
• Non-metals that gain valence electrons more readily are more reactive.
• Lithium requires more energy to lose a valence electron compared to sodium, which has three electron shells versus lithium's two.

Metals Reactivity

• Down a group, reactivity increases; the energy required to remove a valence electron decreases due to increased electron shells, placing valence electrons further from the positive nucleus.
• Across a period, reactivity decreases; the energy required to remove a valence electron increases due to the negative valence electron being closer to the positive nucleus.

Non Metals Reactivity

• Reactivity decreases down a group: having more electron shells makes it harder to attract more electron
• Across a period, reactivity increases due to increased nuclear charge, which pulls valence electrons closer and makes attracting additional electrons easier.

Mendeleev’s Original Periodic Table vs Modern Periodic Table

Similarities

• Both tables arrange elements by increasing atomic properties and group elements with similar properties into columns.
• Many element families (e.g., alkali metals, halogens) remain consistent across both tables.

Differences

• Arrangement basis: Mendeleev’s table used atomic mass; the modern table uses atomic number (proton count).
• Mendeleev's table had gaps for undiscovered elements, with predictions of their properties, while the modern table includes these elements.
• Noble gases were absent in Mendeleev’s table due to their later discovery.
• Element placement: Adjustments in the modern table (e.g., tellurium and iodine) fit atomic number trends rather than atomic mass.
• Lanthanides & Actinides: The modern table includes these as separate rows, absent in Mendeleev’s version.

Group Characteristics

Group 1: Alkali Metals

• These metals are highly reactive, especially with water.
• They form +1 cations (e.g., Na⁺, K⁺).
• They feature low melting and boiling points (compared to other metals).
• They are good conductors of electricity.
• They are soft and can be cut with a knife.

Group 2: Alkaline Earth Metals

• They are less reactive than alkali metals but still highly reactive.
• They form +2 cations (e.g., Mg²⁺, Ca²⁺).
• They display higher melting and boiling points than Group 1.
• They are commonly found in minerals.

Transition Metals

• They exhibit variable oxidation states (e.g., Fe²⁺, Fe³⁺).
• They are good conductors of heat and electricity.
• They have high melting and boiling points.
• They often form colored compounds.
• They are strong and dense.

Post-Transition Metals

• These metals are softer and have lower melting points than transition metals.
• They are poorer conductors of electricity than transition metals.
• They tend to form covalent compounds.

Metalloids

• Show properties intermediate between metals and nonmetals.
• Demonstrate semiconductor behavior (e.g., silicon in electronics).
• Exhibit variable conductivity depending on conditions.

Group 17: Halogens

• Highly reactive nonmetals.
• Form -1 anions (e.g., Cl⁻, F⁻).
• Exist in different physical states at room temperature (F and Cl as gases, Br as liquid, I and At as solids).
• Are strong oxidizing agents.

Group 18: Noble Gases

• Have extremely low reactivity because of a full valence shell.
• Colorless, odorless gases.
• Low boiling and melting points.
• Used in creating lighting and inert environments.

Lanthanides

• Known as rare-earth metals.
• High reactivity similar to alkaline earth metals.
• Used in magnets, electronics, and catalysts.
• Often form +3 cations (e.g., Ce³⁺, Nd³⁺).

Actinides

• Mostly radioactive.
• Include uranium and plutonium, which are used in nuclear energy.
• Form multiple oxidation states.
• High density and high reactivity.

Hydrogen in Group 1

• Hydrogen is in Group 1 because it has one valence electron and forms +1 ions like alkali metals, but hydrogen is a non-metal gas, lacks metallic properties, and mainly forms covalent bonds instead of ionic bonds.

Chemical Bonds

• Atoms form chemical bonds to achieve valence (outermost) shell stability.
• "Octet Rule": Atoms tend to achieve eight electrons in their valence shell.
• Metals tend to lose electrons.
• Non-metals tend to gain electrons.

Definitions

• Soluble: Able to dissolve.
• Insoluble: Does not dissolve.
• Solute: The minor component of a solution (dissolved into the solvent).
• Solvent: The substance that dissolves the solute to form a solution.
• Solution: Mixture of a solute dissolved in a solvent.
• Mixture: Substance containing two or more pure substances physically mixed together.

Concentration

• Concentration can be measured in mol/L or g/L (solute per volume of solution); a larger number indicates a more concentrated solution.
• Dilute: Has less dissolved solute (e.g., 3 mol/L).
• M = n/v (M: Molar Concentration, n: Moles of solute, v: Liters of Solution)

Notes

• Oppositely charged particles attract.
• Salts (ionic substances) are composed of positive and negative ions.
• Water has a slightly positive and slightly negative side.
• The positive side attracts the negative ion and the negative side attracts the positive ion.
• δ (Delta) in chemistry means small (e.g., small positive charge of hydrogen)

Dissolution Equation

• In solution, ions separate and move independently from one another.
• NaCl(s) --> Na+(aq) + Cl-(aq)

Types of Chemical bonds

• Chemical bonding occurs due to the rearrangement of electrons.
• The three types of chemical bonds are metallic, ionic, and covalent.
  • Metallic (e.g. Copper, iron, potassium)
  • Ionic (e.g. sodium chloride)
  • Covalent (e.g. oxygen gas, nitrogen gas)

Properties of Ionic Compounds

• Ionic compounds: Formed between a metal and a non-metal through transfer of electrons, or between a positive and a negative ion.
• Defined as the electrostatic force of attraction between a cation and an anion.
• Polyatomic ions, which involve more than one atom forming an ion, can also form ionic bonds (e.g., ammonium nitrate).
• Ionic substances form a continuous network (lattice).

Conductivity of Ionic Substances

• Some ionic substances are soluble in water.
• Solid State: Cannot conduct electricity as charges (ions) are in fixed positions.
• Molten (liquid)/Aqueous State: Can conduct electricity as charges (ions) are free to move.
• Higher concentration of ionic salt in an aqueous solution increases conductivity.

Melting Point (M.P.)

• M.P.: The energy needed to turn a solid into a liquid.
• Ionic substances have a high melting point that varies based on the strength of attraction; higher charged ions generally have stronger attractions.
• High m.p. signifies strong bonds and that much energy is needed to break or overcome bonds.

Polyatomic Ions

Everyday Examples

Naming Ionic Compounds

• Use the format: metal non-metalide (e.g., NaCl - Sodium chloride)
• Note: Cation comes before anion.

Naming Transition Metal Compounds

• Transition metals with more than one valency require specifying the valency with roman numerals in brackets (e.g., CuO: Copper (II) oxide).

Naming Polyatomic Ion Compounds

• Named the same way as ionic compounds, but use brackets if there is more than one polyatomic ion.
• E.g. Magnesium nitrate = Mg(NO3)2

Writing Formula of Ionic Compounds

• Ions have charges, but ionic compounds do not (they are neutrally charged)
• Can be achieved through understanding and quick maths, or the 'swap and drop' method.
  • Quick Maths: For Magnesium Chloride
    • Mg2+ Cl-
    • 2 + ?(-1) = 0, solve for ?=2
    • Therefore MgCl2

Representing Ionic Substances

• The purpose of the electron dot diagram (aka Lewis structures) is to show how the bonding works between the two ions.
• E.g. NaCl
• Sodium loses an electron to become a cation, and chlorine gains an electron to become an anion.

Necessary Inclusions

• Square brackets around each ion
• Charge in the top right corner
• Valence electrons for the anion only
• If there is more than one ion, use a coefficient to demonstrate how many ions.