Families & Trends in the Periodic Table

Families & Trendsin the Periodic Table Unit 5 Chapter 4

Groups/Families • Alkali Metals • Alkaline Earth Metals • Halogens • Noble Gases • Transition Metals • Inner Transition Metals • Lanthanoids (Rare Earths) • Actinoids

The Little Element That Could • Hydrogen is typically placed in its own category. • Most of the time, Hydrogen behaves like a nonmetal • It can, under certain circumstances, behave like a metal. • Jupiter has a very strong magnetic field because of metallic hydrogen >20 million psi

Representative Elements • Exhibit nearly perfect periodicity. • All members of these groups behave as expected. • Groups on the outside of the table: • Alkali Metals (Group 1 ) • Alkaline Earth Metals (Group 2 ) • Halogens (Group 17 ) • Noble Gases (Group 18 )

Alkali Metals • Group 1 (excluding hydrogen) [ns1] • Soft, lustrous, oxidize when exposed to air. • Difficult to isolate – never found in nature. • React (violently) with water to form a base & H2 gas. • React with chlorine to form a salt with a 1-to-1 ratio: • LiCl • NaCl • KCl • RbCl • CsCl (also FrCl)

Alkaline Earth Metals • Group 2 [ns2] • Harder & more dense than Alkali Metals. • Lustrous, oxidize slowly when exposed to air. • React with water or steam to form a base & H2 gas. • React with chlorine to form a salt with a 1-to-2 ratio: • BeCl2 • MgCl2 • CaCl2 • SrCl2 • BaCl2 • RaCl2 Helium belongs in this group, but you’ll never see: HeCl2

Halogens • Group 17 [np5] • They are all nonmetals • Gases (F, Cl), liquid (Br), and solids (I, At) • Name means “ salt former .” • React with sodium to form a salt with a 1-to-1 ratio: • NaF • NaCl • NaBr • NaI • NaAt

Noble Gases • Group 18 [np6] (except He) • Unreactive Gases – colorless & odorless. • Some of the last natural elements to be discovered. • Once called “Inert Gases” and are all Monatomic (one atom) • When they react, they behave like nonmetals. • They do not combine with other atoms because their outer p-orbitals are full.

Non-representatives • Other families have similarities, but do not behave exactly as expected • Groups 13-16, start with Boron – Oxygen • Usually more differences than similarities • Others are lumped together for other reasons • Transition Metals • Lanthanoids • Actinoids

Transition Metals • Groups 3 to 12 [ndx] • They form the central portion of the Periodic Table. • Behavior and appearance vary. • They have variable oxidation states (charges). • Different oxidation states can produce different colors. • Often used to make pigments. Co+2 Cr+6 Fe+3 Ni+2 Cu+2 Mn+7

Lanthanoids • 1st Row on Bottom of table [4fx] • AKA Lanthanides & Rare Earths • Not so rare (Ce 25th most abundant) • So similar that they are very difficult to separate – remember Moseley? • Most deflect UV – used in sunglasses • Shiny, silvery white, soft, react violently with most nonmetals, tarnish in air

Actinoids • 2nd Row on Bottom of table [5fx] • AKA Actinides • All are radioactive • Not as similar as the Lanthanoids • Only Th and U are common in nature • Most are man-made • Nuclear fallout • Particle colliders

State of the Union • Reacted State: • When elements are combined with other elements to form compounds • Oxidation state ≠ 0 • Most common state (Metals +, Nonmetals usually -) • Elemental State: • When elements are uncombined • Oxidation state = 0 • Most elements are Monatomic (one atom) • Some are always Diatomic (two atoms) • A few are Polyatomic (>2 atoms)

Diatomics • Seven elements always form diatomic molecules when they are isolated in their elemental state…ALWAYS! • Diatomic means “two atoms”. • Two atoms of the element bond together. • Hydrogen, Nitrogen, Oxygen, Fluorine, Chlorine, Bromine, & Iodine • 5 gases, 1 liquid, & 1 solid. • Luckily, Mr. Brinclhof is here to help! You MUST remember All 7 of these! Br2 I2 N2 Cl2 H2 O2 F2

H Another Way • The rule of “7” • Find element #7 • Diatomic elements form a “7” on the Periodic Table (excluding H) 2 2 2 2 2 2 2

The Oddballs The Polyatomic elements: • Sulfur is normally found as S8 • Selenium also forms Se8 • Phosphorus forms P4 Remember! This is for the Elemental State ONLY!

An Allotrope is • When a pure element can be found in more than one physical form • Several elements have different allotropes, but most often cited is Carbon • Carbon has 3 common allotropes • Amorphous – Random arrangement of C atoms • Graphite – Hexagonal arrangement in sheets • Conducts electricity along one axis only! • One single sheet is called Graphene. • Diamond – 3-D network solid

Allotropes of C Graphene Amorphous C Graphite (Sheets) Diamond (Network Solid)

Trends in the Periodic Table • Several trends appear once we have the elements in order on the Periodic Table • Atomic Radius • Electronegativity • Ionization Energy • Reactivity

Ray “D” Us • The atomic radius decreases from left to right • It increases from top to bottom Na is bigger than Ar (223 pm) (88 pm) Na Mg Al Si P S Cl Ar F I is bigger than F (132 pm) (57 pm) Cl Br I

Fluorine says, “Mine!” • Electronegativity is a measure of how badly an element wants to gain an electron • It increases from left to right • It decreases from top to bottom

F has Codependency Issues • Ionization Energy is the amount of energy required to remove an electron. • It increases from left to right • It decreases from top to bottom Na needs less NRG than Cl (496 kJ/mol) (1256 kJ/mol) F needs more NRG than I (1681 kJ/mol) (1008 kJ/mol)

Major Trends in a Nutshell Fluorine has the smallest radius, highest electronegativity, and highest ionization energy! (excluding the noble gases) F Francium has the largest radius, lowest electronegativity, and lowest ionization energy! Atomic Radius Decreases Electronegativity Increases Ionization Energy Increases Fr We usually ignore the Noble Gases simply because they do not react.

The Reactivity trend depends on the type of element… • The most reactive Metals are farther down and to the left • The most reactive Nonmetals are higher and to the right Fr F

Why, why, why!?! • The periodic properties of elements stem from the repeating patterns of electrons. • The varying trends across the periodic table stem from each electron’s attraction to the nucleus. • Since electrons repel each other, the more electrons between the outer electron and the nucleus, the less “pull” the electron feels. • This effect is called shielding.

Z F - what?!? • The effective charge that an electron feels is: • Zeff = Z – σ • Where Z is the nuclear charge (# of protons) and • σ is a measure of how many other electrons are between a specific electron and the nucleus. • The lower the effective charge, Zeff, the farther out an electron will “wander.” • Wandering electrons increase the size of the atom & the likelihood that the atom will lose that electron. • John Slater came up with a set of rules that approximate the value of σ in 1930

Exceptions to Configurations • When we look at the electron configurations of the elements, we see some exceptions: • Orbitals tend to be stable when they are • 1. Half Full • 2. Completely Full • A greater effect is seen on the stability of more complex orbitals. • Because of this, certain elements have electrons in the “wrong” orbital.

The Rebellion begins! • Chromium should be • 1s22s22p63s23p64s23d4 • But when we check experimentally, we find: • 1s22s22p63s23p64s13d5 • Draw the box diagram for 4s and 3d • 3d is more stable, energetically, half full 4s 3d

Why would you do this!?! • As the electrons get farther away from the nucleus, • The closer the energy levels get to each other (look at your H-Spec diagrams – you only drew the s-orbitals). • The difference between 4s and 3d is fairly small. • When 3d is half full (completely full, too!), it actually drops in energy (on par with 4s) • 1 electron is promoted (2 for Pd & Th).

The Rebels • This only happens in the d- and f-blocks. • The d-block perpetrators are: • Pd 4: Cr (Ni) Cu • Pd 5: Nb Mo Ru Rh Pd Ag (Pd has no 5s!) • Pd 6: Pt Au • The f-block has: (note: Lu & Lr are d-block) • 4f: La, Ce, & Gd (La is [Xe]6s25d1 – no 4f!) • 5f: Ac, Th, Pa, U, Np, & Cm (Ac & Th have no 5f!)

Families & Trends in the Periodic Table