Understanding the Periodic Table

Periodic Table Basics

• A group is a column in the periodic table; there are 18 groups total;
• A period is a horizontal row, which indicates the number of electron shells an element has.

Element Types

• Metals are on the left side of the periodic table; they are malleable, lustrous, have high melting points and thermal conductivity, lose electrons to form positive ions.
• Non-metals are on the right side of the periodic table; they are brittle, dull, have lower melting points and thermal/electrical conductivity, and gain electrons.
• Metalloids have intermediate properties between metals and non-metals.
• Electron configuration can be easily predicted; elements in a group have a common number of valence electrons, and the period number indicates the outer energy level.

Block Method

• The block method is used for electron configuration to identify elements.
• Use the period number to determine the outermost energy level.
• Write the electron configuration as you move along each row.

Specific Groups

• Alkali metals (Group 1) tend to form alkaline solutions when reacted with water, are soft metals that are reactive with water and many non-metals.
• Transition elements show transitionary behavior between s and p block elements and can form ions with different ionic charges.
• Halogens (Group 17) means "salt producing," are very reactive, and form ionic salts when reacted with metals.
• Noble gases (Group 18) are colorless, monoatomic, and have little to no reactivity with other elements.

Atomic Properties

• Atomic radius is half the distance between two adjacent nuclei.
• Effective nuclear charge is the net positive charge experienced by an outer valence electron, resulting from attraction between protons and repulsion between inner core electrons.
• As you move across a period, the effective nuclear charge increases.

Atomic Radii Trends

• Atomic radii decrease across a period due to the greater effective nuclear charge.
• Atomic radii increase down a group because elements are larger.
• Ionic radii for positive ions are smaller than their parent atom due to the loss of valence electrons and increased attraction between the nucleus and remaining electrons.
• Greater positive charge means more electrons removed and smaller radius.
• Negative ions are larger than their parent atom due to increased repulsion when electrons are gained.
• More negative charge means more electrons gained and increased radius.
• Isoelectronic species have the same electron configuration and compare radii.

Atomic Radius

• Atomic radius decreases across a period because valence electrons are in the same energy level, but increasing nuclear charge pulls electrons more strongly.
• Atomic radius increases down a group because additional energy levels are added.

Ionization

• Metals tend to lose electrons; non-metals tend to gain electrons.
• Electron affinity is the amount of energy released when a neutral atom gains an electron to form a negatively charged ion.
• First electron affinity occurs when a neutral, gaseous atom gains one electron, releasing energy due to attraction between the nucleus and incoming electron.
• Down a group, the absolute value for first electron affinity decreases (lower affinity for gaining an electron), and elements further down have larger ionic radii with outer valence shells away from the nucleus.
• Weaker attraction between the added electron and the nucleus cause an increased absolute value (* left pointing arrow *) for 1st electron affinity going across a period (higher affinity for gaining an electron) - across the period: more effective nuclear charge increases, resulting stronger attraction between added e- and nucleus.

Ionization Energy

• Ionization energy is the energy required to remove the outermost electron from a species.
• First ionization energy is the minimum energy needed to remove one mole of electrons from one mole of gaseous atoms at the ground state.
• Down a group, first ionization energy decreases because atomic radii increase, electrons occupy higher energy levels, and there is weaker attraction to the nucleus.
• Less energy is required to remove an electron when moving down a group.
• Across a period, first ionization energy increases because nuclear charge increases (more protons in the nucleus).
• Atomic radius decreases (except Ar) accross a period, so more energy is required to remove an electron.

Electronegativity Trends

• Electronegativity is the tendency of an atom to attract a shared pair of electrons in a covalent bond.
• Electronegativity decreases down a group as atomic radius increases from additional levels, as the distance between nucleus increases and the shared pair of electrons; weaker attraction.
• Electronegativity increases across a period because effective nuclear charge increases, there is a stronger attraction between the nucleus and the shared pair of electrons.
• Noble gases don't have an electronegativity value because they don't form shared pairs of electrons with other atoms.

Metallic Character

• Metallic character is the tendency of an element to lose electrons to form positive ions, and increases as you move to the left and down on chart.
• Lower ionization energy means a greater tendency for electron loss and greater metallic character.

Non-Metallic Character

• Non-metallic character is the tendency of an element to gain electrons to form negative ions, and increases as you move up and to the right on chart.
• Higher electron affinity means a greater tendency for electron gain and greater non-metallic character.

Group 1 Metals

• Group 1 metals are highly reactive (e.g., with water) and stored in oil.
• They react with water to produce alkaline solutions and hydrogen gas.
• Reactivity increases as we move down group one, as reactivity increases L-->R they are losing their valence electrons; the distance increases between the valence electrons down a group, so we need higher ionization energy.

Oxides, Halides, & Ocean Acidification

• Oxides of metals are basic, nonmetals are acids
• Halides, group 17: Reactivity decreases down group
• Ocean acidification: Caused by <accumulation of CO2> in atmosphere (combustion of fossil fuels): Dissolving CO2 in water increases acidity = carbonic acid is formed: bad for marine animal food chains!

Acids and Bases

• Acids have a pH between 0 and 7 (battery acid, vinegar).
• Bases have a pH between 7 and 14 (blood, bleach).
• Amphoteric substances behave as either an acid or a base.
• Metallic and non-metallic oxides

Oxidation State

• Oxidation state is a number assigned to an element showing the number of electrons lost, gained, or shared in a compound or ion, based on relative electronegativity, and are hypothetical.
• A pure element oxidation state =0, distribution/election is the same.
• Monoatomic ions have an oxidation state = to charge of the ion and its charge.
• Oxidation state of all the atoms in a neutral compound (=ionic, molecular or both) adds up to zero OVERALL.
• Assigned to give the overall (charge of ion).
• The oxidation states of atoms = the charge of the more common ion
• Assign positive #s to the more metallic elements first and so on to non-metallic elements first.