Kaplan Organic Chem Chapter 2: The Periodic Table

The periodic table arranges the elements into

• periods(rows) and groups/families(columns) based on atomic number

• groups have elements that have the same electronic configuration in their valence shell and share similar chemical properties

Valence electrons are the

• farthest from the nucleus and have the greatest amt of PE

• valence shell electrons largely determine the chemical reactivity and properties of the element

What are A and B elements?

• A are representative elements and have their valence electrons in s or p subshells

• B are nonrepresentative elements and include both transition elements(valence e in s and d subshells) and lanthanide/actinide series(valence e in s and f subshells)

• for the representatitive elements, the roman numeral and letter designation determine the electron configuration: ex→element in group VA has 5 valence eletrons with s²p³ configuration

What are the physical and chemical properties for metals?

• shiny solids, except for mercury which is liquid under standard conditions

• high melting points and densities, except for lithium which has a density half of water

• ability to be deformed without breaking called malleability

• ability to be pulled or drawn into wires called ductility

• low effective nuclear charge, low electronegativity(high electropositivity), large atomic radius, small ionic radius, low ionization energy, and low electron affinity

• can easily give up electrons

Many of the transition metals have 2 or more oxidation states(charges when forming bonds with other atoms). Because the valence electrons are only loosely held to their atoms, they are

• free to move, which makes metals good conductors of heat and electricity

Valence electrons of active metals are found in the

• s subshell

• those of transition metals found in s and d subshells

• those of lanthanide and actinide series in s and f subshells

What are the physical and chemical properties of nonmetals?

• brittle in solid state ad show little to no metallic luster

• high ionization energies, high electron affinities, high electronegativity, small atomic radii and large ionic radii

• poor conductors of heat and electricity

• can’t give up electrons easily

What are the physical and chemical properties of metalloids/semimetals?

• electronegativities and ionization energies lie between those of metals and nonmetals

• everything else varies; reactivities are dependent on the elements with which they are reacting

What is the effective nuclear charge(Zeff)?

• if you move from left to right on PT→electrons and protons added one at a time→as positivity of nucleus inc→electrons surrounding the nucleus experience a strong electrostatic pull toward the center of the atom→causes electron cloud to move closer and bind more tightly to nuc→electrostatic attraction btwn the valence shell electrons and nucleus id Zeff, a measure of the net positive charge experienced by the outermost electrons

• for elements in the same period, Zeff increases from left to right

As one moves down the elements of a given group, the principal quantum number

• increases by one→this means that the valence electrons are increasingly separated from the nucleus by a greater number of filled principal energy levels(inner shells) →result is reduced electrostatic attraction btwn the valence electrons and pos charged nuc

• these outermost electrons are held less tightly as the principal quantum number increases

• as one goes down in a group, the increased shielding created by the inner shell electrons cancels the increased positivity of the nucleus

• the valence electrons are held less tightly to the nucleus as one moves down a group due to the increased sep btwn valence electrons and the nuc

Elements tend to be most stable with

• 8 electrons in their valence shell

What is the periodic table trends?

• Zeff: increases left to right; this is bc electrons being added only to outermost shell and inner shell electrons remain constant→increasing pos charge of the nuc pulls the outer electrons more closely inward and holds them more tightly; the greater the atom’s Zeff, or the closer the valence electrons are to the nuc, the more tightly bound they are

• Atomic radius: increases from right to left and down; increasing principal quantum num→valence electrons found farther away from nuc bc number of inner shells increasing

• Ionic radius: for nonmetals close to metalloid line, their group number says they require more electrons than other nonmetals to achieve stable octet→these gain electrons while their nuclei maintain the same charge→have a larger ionic radius than those closer to 18;

• for metals close to metalloid line, they have more electrons to lose to achieve stable octet→smaller ionic radius than other metals; metals closer to group 1 have fewer electrons to lose and experience a less drastic reduction in radius

• Ionization energy: increases to the right and up; this is the energy required to remove an electron from a gaseous species, which requires heat, making it an endothermic process; the higher the Zeff, the higher the IE

• electron affinity:increases to the right and up; this is the energy let off by a gaseous species when it gains an electron, which releases heat, making it an exothermic process; the strong the electrostatic pull(higher Zeff) btwn nuc and valence e, the greater the energy release will be when the atom gains the e; bc the valence shell is farther away from the nuc as the principal quantum number inc, electron affinity dec in a group from top to bottom; group 18 have 0 EA bc they already have a stable octet

• Electronegativity: increases to the right and up; a measure of the attractive force that an atom will exert on an electron in a chemical bond; first 3 noble gases have high IE, but no EN

What is first and second ionization energy?

• subsequent removal of a second or third electron requires even more energy bc the removal of more than 1 means the electrons are being removed from an increasingly cationic positive species

• elements in groups 1 and 2 have such low IEs that they’re called active metals→these do not exist naturally in their neutral forms and are always found in ionic compounds; they will lose electrons to achieve stable oct

• group 17 elements don’t give up their electrons, and in their ionic form, they are generally anions

Describe the chemistry of alkali metals(group 1).

• densities lower than other metals

• react readily with nonmetals, esp halogens; react violently with water

• alkali and alkaline earth metals are called the active metals bc they are so reactive that they’re not found in their elemental (neutral) state

Describe the chemistry of chalcogens(group 16).

• at high concentrations, many of these can be toxic/damaging

Describe the chemistry of halogens(group 17).

• these are so reactive that they are not naturally found in their elemental state but rather as ions or diatomic molecules

Describe the chemistry of noble gases.

• minimal chemical reactivity due to their filled valence shells

• low boiling points and exist as gases at room temp

The transition elements(groups 3 to 12) are considered to be metals and have

• low EA, low IE, and low EN

• very hard and have high melting and boiling points

• malleable and good conductors due to the loosely held electrons that fill the d orbitals in their valence shells

• many of them can have diff possible charged form or oxidation states(different vibrant colors) bc they are capable of losing diff numbers of electrons from the s and d orbitals in their valence shells→bc of this, these metals can form many diff ionic compounds

When we perceive an object as a certain color, it is bc

• that color is not absorbed, but rather reflected by the object

• if an object absorbs a given color of light and reflects all others, our brain mixes these subtraction frequencies and we perceive the complementary color of the freq that was absorbed

• ex: carotene strongly absorbs blue light but reflects other colors→out brain interprets the color of carotene as the result of white light minus blue light, which is yellow light

Which has the larger radius: F or F-? K or K+?

• F-, K →the ionic radii of anions are larger, but ionic radii of cations are smaller