The Periodic Table and Electron Configuration

Periodic Table

• Elements are arranged in order of increasing atomic number.
• The periodic table consists of metals, metalloids, and non-metals.
• It includes 7 horizontal rows called periods.
• It includes 18 vertical columns called groups.
• The arrangement of elements relates to their electron configuration.
• Group 1 elements are Alkali Metals
• Group 2 elements are Alkaline Earth Metals
• Group 3-11 elements are Transition Metals
• Group 17 elements are Halogens
• Group 18 elements are Noble Gases.
• Atomic number indicates the number of protons in an atom.
• Relative atomic mass, located under the atomic symbol, is relative to carbon-12.

Electron Configuration

• Electron configuration is the distribution of electrons around an atom's nucleus.
• A carbon atom with 6 electrons has 2 electrons in the first shell and 4 in the second (valence) shell
• Carbon's electron configuration is 2.4.
• Repulsion between electrons limits the number in a shell.
• Maximum number of electrons per shell:
  • 1st shell: 2
  • 2nd shell: 8
  • 3rd shell: 18
  • 4th shell: 32
• Calcium (Ca) has an electron configuration of 2.8.8.2.
• Calcium has four electron shells and two valence electrons.
• It's in Period 4 and Group 2.

Trends in the Periodic Table

• Elements in the same period have the same number of occupied electron shells, while proton and valence electron numbers increase.
• Down a group, elements share the same number of valence electrons, but the number of electron shells increases.

Atomic Radius

• Atomic radius signifies the distance from the nucleus center to outermost electron shell.
• Atomic radius generally increases down a group.
• Elements down a group have the same number of valence electrons, but more electron shells.
• Increased electron shells shield valence electrons, reducing attraction to the nucleus and increasing atomic radius.
• Atomic radius generally decreases across a period.
• As protons increase across a period, nuclear charge becomes more positive.
• Outermost electrons experience similar shielding, leading to stronger attraction to the nucleus and reduced atomic radius.

Ionic Radius

• Ionic radius is the distance between a nucleus and the electron in the outermost shell of an ion.
• Ionic radius increases down a group, similar to atomic radius.
• Ionic radius decreases across a period for ions of the same type.
• Ions with the same electron number and configuration are isoelectronic.
• A cation is smaller than its neutral atom while an anion is bigger.

Ionization Energy

• First ionization energy is the energy needed to remove one mole of outermost electrons from one mole of gaseous atoms to form 1+ ions.
• First ionization energy typically increases across a period as it becomes harder to remove an electron.
• Ionization energy increases because of increased positive nuclear charge and a smaller atomic radius.
• First ionization energy typically decreases down a group as electrons are easier to remove.
• Outermost electrons are further from the nucleus with more shielding, resulting in a weak attraction.

Electronegativity

• Electronegativity indicates an atom's ability to attract electrons in a chemical bond.
• High electronegativity atoms attract electrons more strongly.
• Low electronegativity atoms tend to lose their own electrons.
• Electronegativity decreases down a group.
• Moving down a group, atoms are larger which keeps bonding electrons further away from the nucleus. Thus, atoms further down the group have less attraction for bonding electrons
• Electronegativity increases across a period because atoms have greater nuclear charge and a smaller covalent radius which allow the nucleus to attract bonding electrons more strongly.

Group 2 Elements

• Known as alkaline earth metals, they are soft, silver metals that are harder and denser than Group 1 metals.
• They become more reactive down the group.
• These elements form +2 charged cations via loss of two electrons.
• As reducing agents, Group 2 metals form cations.
• Barium is most easily oxidized, making it the strongest reducing agent, with Beryllium as the weakest.
• Group 2 elements react with oxygen to form oxides.
• They react with halogens to form halides.
• Magnesium burns brightly in steam, producing magnesium oxide and hydrogen gas.
• Magnesium reacts with warm water to produce hydroxide and hydrogen gas.
• Other Group 2 elements (excluding beryllium and magnesium) react with cold water form metal hydroxides and hydrogen gas.
• Group 2 hydroxides' solubility increases going down the group and result is alkaline water.
• Magnesium hydroxide is insoluble in water, but Calcium hydroxide is reasonably soluble.
• Group 2 carbonates decompose with heating to produce oxides oxides and carbon dioxide gas.
• Thermal stability of Group 2 carbonates increases down the group.
• Metal ions are bigger down Group 2 but have the same charge which means the charge density is reduced going down the group.
• A metal cation with a low charge density has weak polarising power, meaning the carbonate ion is less polarised, and therefore more thermally stable.

Group 17 Elements

• Known as halogens which exist as diatomic molecules:
  • F2 (gas)
  • Cl2 (gas)
  • Br2 (liquid)
  • I2 (solid)
• Reactivity decreases down this group.
• They gain one electron to form anions with a -1 charge.
• Melting & boiling points of the halogens increase down the group.
• Fluorine, chlorine, bromine, and iodine are non-polar molecules.
• Weaker Van der Waals forces exist between the molecules.
• As molecules get larger and have more electrons, temporary dipoles can be set up
• Stronger Van der Waals forces exist between larger molecules.
• Group 17 elements are oxidising agents.
• Fluorine has the highest tendency to be reduced, and is the strongest oxidising agent. Iodine is the weakest oxidising agent.
• A strong oxidizing halogen displaces those with lower power from compounds via displacement reactions.
• Halogens react with hydrogen but at varied rates that depend per halogen.
• Halogens react with water, forming halides and hypohalides at varies rates.
• Fluorine displaces oxygen gas and makes hydrogen fluoride.

Reactions of Halide Ions

• Halide ions react with silver nitrate to indicate which halide ion is present.
  • Fluorine has no reaction resulting in no precipitate
  • Chlorine results in a white precipitate
  • Bromine results in a cream participate
  • Iodine results in a pale yellow precipitate

Reactions of Silver Halide Precipitates with Ammonia

• Silver halide precipitates are treated with ammonia to help differentiation.
  • Silver chloride dissolves in dilute ammonia, forms colourless complex ion
  • Silver bromide dissolves in concentrated ammonia, forms colorless complex ion
  • Silver Iodide will not react with Ammonia because the reaction is insoluble.

Hydrogen Halides

• Hydrogen fluoride and Hydrogen chloride are very stable to heat and do not split up into hydrogen again with heat.
• Hydrogen bromide splits slightly.
• Hydrogen iodide splits to a greater extent.
• Strength of the hydrogen-halogen bonds decreases down the group.
• As Group 17 atoms are getting bigger down the group, the bond become longer and weaker.
• Group 17 halides are less thermally stable, because weaker bonds need less heat to break.

Explore the periodic table's organization by atomic number, metals, and groups. Learn about electron configuration, including distribution around the nucleus and valence shells. Understand electron capacity in shells and carbon's electron configuration.