The Periodic Table

The Chemistry of Groups

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Noble Gases

Elements found in group VIII (Group O) They are called the inert gases -They do not make an chemical combinations with any of elements -There aren't any compounds with this group -All gases at room temperature -Very low boiling and melting point -They put a color in the in visible wavelengths when low pressure of a gas is put into a tube and high voltage is run through it, called a neon light -Fairly nonreactive

Halogens

Elements in Group VIIA -highly reactive nonmetals with 7 valence electrons, one short of octet -highly favorable physical properties -They range from gaseous (F2 and Cl2), Liquid (Br2), Solid (I2) at room temp -More uniform chemical properties -High electronegativity -reactive towards alkali metals and alkaline earths, they want grouse that will readily donate electrons to form stable ionic crystal. -Fluourine (F) has the highest electronegativity.

Periodic Law

States that the chemical properties of the elements are dependent on their atomic number

Electron Affinity

this is the energy change that occurs when an electron is added to a gaseous atom and represents how easily an atom can accept an electron. -The stronger the attractive pull of nucleus for electrons (Effective nuclear charge/Zeff) the greater the electron affinity will be. -A positive electron affinity value represents energy release when an electron is added to an atom.(Commonly used) -A negative electron affinity represents a release of energy. (Not used)

Periodic properties of the elements

All elements seek to gain or lose electrons as to gain a stable octet configuration (Which is possessed by the inert or noble gasses of group 8.) There are two periodic trends that exist when looking at the periodic table. -1st: as one goes across from left to right on the periodic table, electrons are added one by one. The electrons of the outermost shell experience an increasing amount of nuclear attraction, becoming closer and more tightly bound to the nucleus. -2nd: As one goes down a given column, the outermost electrons become less tightly bound to the nucleus. This is because the number of filled principal energy levels increases downward within each group.

Transition elements

Elements found in Group IB to VIIIB -All are considered metals and are called the transition metals -Very hard and have high BP and MP -As you move across a period, the 5 d orbitals become more filled -D electrons are held loosely by the nucleus and are mobile -The mobility of the electrons contributes to the malleability and high electrical conductivity -Low ionization energies -Can exist in a variety of positively charged forms or oxidation states. -Capable of losing various electrons from the S and D orbitals of their valence -They COULD have oxidation states from +1 to +8, but they don't exhibit so many -The ability of the transitions to have positive oxidation states allows them to have many different ionic and partially ionic compounds. Complex ions form with dissolved ions and can form with: -water molecules (Hydration complexes) or -Nonmetals-highly colored solutions and compounds (CuSO4 5H20) which can enhance the low solubility of certain compounds Ex. AgCl is insoluble in water, but is soluble in aqueous ammonia because of the complex that forms [Ag(NH3)2]+ -Formation of complexes causes the d orbitals to split into two energy sublevels and enables the absorption of certain frequencies of light by the complexes (The frequencies containing the precise amount of energy to raise electrons from the lower to higher d sub level. -Subtraction frequencies- are the frequencies not absorbed and give the complexes their characteristics Ex. Copper (Cu) group IB can exist in the +1 or +2 oxidation states Ex 2. Manganese (Mn) Group VIIB occurs in the +2,+3,+4,+6,+7 oxidati states

Alkali Metals

Elements of Group IA -Posses most of the qualities of metals, but their densities are lower than other metals. -One loosely bound electron in their outermost shell -Have the largest atomic radii because of the one electron in the outer shell -Low ionization energy (They can easily loose their valence to form a univalent cation) -High reactivity and metallic properties -LOW electronegativities -Highly reactive with nonmetals, esp Halogens

Alkaline Earth Metals

Elements of Group IIA -Possess metallic properties -Their properties depends on their ease of losing electrons (They have two electrons in their valence) -Smaller atomic radii than the Alkali -Two electrons are easily removed to form a divalent cation -Low electronegativity -Positive electron affinity (because it's very difficult for them to gain 6 more electrons than for them to lose the 2)

Metalloids (Semimetals)

Found on the line between the metals and the nonmetals in the periodic table Nonmetals are: Boron (B), Silicon (Si), Germanium (Ge), Arsenic (As), Antimony (Sb), and Tellurium (Te). -Properties vary considerably -Densities, melting points, and boiling points fluctuate considerably. -Electronegativites and ionization energies lie between the metals and nonmetals (They contain characteristics of both classes) -Boron (B) is more likely to form covalent bonds like a nonmetal than to donate them like a metal. -The reactivity of metalloids is dependent on what they are reacting with. (See Ex. 2) -Metalloids are semi-conductors- they have the ability to conduct electricity, but not to full conduction. Ex. Silicon has a metallic luster, is brittle, and is not an efficient conductor of electricity. Therefore, it has the metallic luster of metals, but the brittleness and inefficient conductivity of a nonmetal. Ex 2. Boron (B) behaves as a nonmetal when reacting with Sodium (Na) and as a metal when reacting with Fluorine (F).

Electron Affinity Generalizations

Generalizations can be mad about the electron affinities of a particular group in the periodic table -Group IIA elements (Alkaline Earths) have low electron affinities. The reason for this is because they're stable due to the filled s subshell. -Group VIIA elements (Halogens) have high electron affinities because the addition go an electron to atom results in a completely filled shell-->Stable electron configuration Achieving the stable octet involves a release of energy and the strong attraction of the nucleus for the electron leads to a high energy change. -Group VIII elements (Noble Gases) have electron affinities on the order of zero (they have a stable octet) and cannot readily accept an electron. -Elements of other groups generally have low values of electron affinity.

Electronegativity

Is a measure of the attraction an atom has for electrons in a chemical bond. -The greater the electronegativity of an atom, the greater its attraction for bonding electrons. -The Pauling electronegativity scale inhere the values range from 0.7 for the most electropositive (Cesium) to 4.0 for the most electronegative (Fluorine). -Fluorine is the most electronegative element and has the largest (most exothermic) electron affinity. Electronegativities are related to ionization energies: Elements with low ionization energies will have low electronegativities because their nuclei do not attract electrons strongly. -Elements with high ionization energies have high electronegativities because of the strong pull the nucleus has on electrons. Electronegativity increases from left to right across periods. -The electronegativity decreases as the atomix number increases, as a result of increased distance between the valence electrons and the nucleus....greater atomic radius

Nonmetals

Located on the upper right side of Periodic table -They have 4,5,6, or 7 electrons in their outer shell -When they join with other elements, they share electrons in a covalent bond or gain electrons to become a negative ion and form an ionic bond -Nonmetals form covalent bonds with other nonmetals -Are brittle in the solid state and show barely any metallic luster -High ionization energies -High electronegativity -Poor conductors of heat and electricity -Have the ability to gain electors easily -Have a wide range of chemical behaviors and reactivities -Separated by a line cutting diagonally through the region containing partially filled P orbitals. -have lower densities, melting points, and boiling points than that of metals. -Nonmetals are not as cohesive as metals and are not as ductile as metals. -Nonmetals for diatomic or polyatomic ions with other molecules or the same element -Many nonmetals are Allotropes (have different free forms of the same element) that appear under different conditions Radical/polyatomic ions- The formation of groups of elements (usually nonmetals) with covalent bonds that have a common charge. Elemental nonmetals- have a dull appearance, are more likely to be brittle and can shatter when struck

What is the difference between Periods and Groups and how many are there of each?

Periods are the rows going down and the groups are the columns going across. There are 7 periods and 18 groups. -The seven periods represent the principal quantum number from n=1 to n=7. Each period is filled sequentially. -Groups represent elements that have the same electronic configuration in their valence and share similar chemical properties. -There are two sets of groups that take on two designations the A and the B designations. -A group elements are the representative elements and have s and p orbitals in their outermost orbitals. -B group elements are the non representative elements and include transition elements with have partly filled d sublevels, lanthanides, and actinides which have partially filled f sublevels.

Metals

Shiny Solids (Except for mercury (l)) at room temperate and have high melting points and densities. -They have one, two, or three electrons in their outer shell -Metals are more likely to lose electrons and become positive ions, this makes them stable -They attach to other elements to and make ionic bonds -Metals can be deformed without breaking. -Malleability- is the ability for metal to be hammered into shoes. -Ductility-is the ability for metal to be drawn into wires. -Metals form electron gases-when clusters of metals are grouped together with semi-loose electrons floated around the atoms. Electron gases account for the luster of metals -Electron gas also accounts for the conductivity of metals. -Active metals react with acids and some will react with water -Metals are also much more dense than non-metals Metals have the characteristics of having a large atomic radius, low ionization energy, & low electronegativity -The reason for these characteristic is because the few electrons that they have in their valence can be removed. -Metals are good conductors of heat and electricity. -Group IA and IIA represent the most reactive metals. -Transition element are metals which have partially filled d orbitals. Ex. Mercury has a cohesiveness (Allows it to stick to itself and other metals easily) which is a result of electron gas. Ex 2. Silver is very malleable and will likely change shape in its solid shape as opposed to shattering. Ex 3. Aluminum is a metal that is Amphoteric, it reacts with both acids and bases.

Atomic Radii

The atomic radius is equal to one-half the distance between the centers of two atoms of that element that are just touching each other. -The atomic radius decreases across a period from left to right -The atomic radius increases down a a given group. -Atoms with the largest atomic radius will be located at the bottom of the groups and in group 1. -The effective nuclear charge increases across a period and causes the atomic radii to decrease. The reason for this is because the molecules, even though getting more electrons (negative) are also getting more protons, which creates a more positive charge and has more effect on the electrons negative charge.

Types of Elements

There are three main categories of element classifications: Metals (located on the left side and in the middle of the periodic table), Nonmetals (Located on the right side of the table), and metalloids (semimetals) found along a diagonal line between the other two.

Ionization Energy (IE)

This is the energy requires to completely remove an electron from a gaseous atom or ion. -Removing an electron from an atom always requires an input of energy (is endothermic). -The closer and more tightly bound an electron is to the nucleus, the more difficult it will be to remove, and the higher the ionization energy will be. -The first ionization energy is the energy required to remove one valence electron from the parent atom. -The second ionization energy is the energy needed to remove a second valence electron from the univalent ion to form the divalent ion. -Successive ionization energies grow larger as you remove more valence electrons. Ionization energy increases from left to right across a period as the atomic radius decreases. -Moving down a group, the ionization energy decreases as the atomic radius increases. -Group 1 elements have low ionization energies because the loss of an electron results in the formation of a stable octet. -Going from down to up, the electronegativity increases, the ionization energy increases, and the decreasing atomic radius.