PERIODIC TRENDS
ATOMIC SIZE
There is no sharp boundary to an atom, but we can get a rough measure of atomic size from the most probable electron distance from the nucleus (called the atomic radius).
Atomic radius decreases as you move across a period and increases as you move down a group.
IONIZATON ENERGY
This is the energy required to remove an electron from an atom in the gas phase.
A(g) + E ---- A+(g) + e-
For each successive ionization more energy is needed.
Metals generally have low ionization energies, while nonmetals have high ionization energies.
Ionization energy increases as you move across a period and decreases as you go down a group or family.
ELECTRON AFFINITY
This is the energy used or released for a gaseous atom to gain an electron.
A(g) + e- --- A-(g) + E
In general, electron affinity increases (becomes more negative) as you go across a period and decreases as you go down a group or family.
In general, the greater the electron affinity the greater the ionization energy. Metals have lower electron affinities. Nonmetals have higher electron affinities.
ELECTRONEGATIVITY
Both electron affinity and ionization energy deal with isolated atoms. Chemists need a comparative scale relating the abilities of elements to attract electrons when their atoms are combined.
The relative tendency of an atom to attract shared electrons to itself, when bound to another atom, is called electronegativity.
The most active metals (to the left and down on the periodic table) have the lowest electronegativities. Most nonmetals (to the right and up on the periodic table) have the highest electronegativities.
PERIODIC TABLE
1870 – Dmitri Mendeleev
How are electron configurations related to periodic trends?
You need to look at the outer orbit electrons of the elements.
REVIEW OF THE PERIODIC TABLE
Families or groups share chemical properties and have similar electron patterns. They tend to behave alike.
The A groups are the main group elements.
IA – alkali metals
IIA – alkaline earth metals
VIA – chalcogens
VIIA – halogens
VIIIA – inert elements or Noble gases
The B groups are called the transition metals.
IB – VIIIB
IB – coinage metals
The elements in a period or series vary on a regular basis (horizontal rows).
The inner transition metals include the lanthanide and actinide series.
OCTET RULE
Elements with 8 electrons in their outer energy level are particularly stable (2 electrons are stable for hydrogen).
The number of electrons in an atoms outer orbit is equal to the group number for groups IA-VIIIA.
All of the B groups, except for IB (1 outer orbit electron for elements in IB), have 2 outer orbit electrons.
The lanthanide and actinide series also have 2 outer orbit electrons.
METALS, NONMETALS, AND METALLOIDS
Metals are found to the left of the metal-nonmetal line on the periodic table of elements. Metals have only a few electrons (most have 1, 2, or 3) in their outer most energy level. They lose electrons (1, 2, or 3) to form positive ions. They are good conductors, malleable, ductile, etc.
Nonmetals are located to the right of the metal-nonmetal line. They have more than 3 electrons in their outermost orbit and gain electrons to form negative ions.
Metalloids have the properties of metals and nonmetals. They are the elements that are positioned along the metal-nonmetal line.
An example of a metalloid is silicon.
BONDING
Types of Bonds
(1)ionic – attraction between oppositely charged particles caused by a transfer of electrons.
Cu+2 + Cl-1 --- CuCl2
Na+1 + Cl-1 --- NaCl
(2)covalent – a bond formed by sharing one or more pair(s) of electrons between atoms.
H2
CCl4
(3)metallic – Electrons in metallic bonds are freely mobile (delocalized). It takes little energy for atoms to surrender their outermost s and p electrons to a “sea” of free electrons.
Metallic bonds consist of the attraction between a network of positively charged metal ions and the surrounding electron “sea”. This generalized electrostatic attraction to freely mobile electrons is a characteristic of pure metals and most alloys (mixtures of metals).
The valence number of an atom is the number of bonds that an atom actually forms.
Electrons that can participate in any kind of bonding (electrons that can be gained, lost, or shared) are called valence electrons. Valence electrons are located in the outermost orbital or valence shell of an atom.
You can determine the number of outer shell electrons by looking at the group number for the A groups (the number of valence electrons is equal to the group number).
All of the B groups and the lanthanide and actinide series have 2 electrons in their outermost shell.
The only exceptions are elements in group IB, which have only one valence electron.
O – VIA – 6 valence electrons
C – IVA – 4 valence electrons
Na – IA – 1 valence electron
G.N. Lewis (electron dot) diagrams show valence electrons and the possibilities for bonding in an atom.
The atomic symbol stands for the nucleus plus all of the underlying, normally filled orbitals.
Valence electrons are represented individually by dots placed in an imaginary square around the atomic symbol in a pairing pattern set by the aufbau principle and Hund’s Rule.
This means that electron dots representing the lower-energy s orbital electrons are paired first, while dots representing the higher, equal energy p orbital electrons are paired only after each side of the square has one.
Fill with electrons based on the numbers listed above (order of fill).
X = symbol of the element which represents the nucleus of an atom plus all of the inner orbit electrons
A dot is used to represent a valence electron.
ELEMENT - GROUP – VALENCE e-‘s - DIAGRAM
Monatomic Ions
The e-dot diagrams of anions and cations will change depending on the number of electrons gained or lost.
Metals (lose electrons)
Nonmetals (gain electrons)
n is the number of electrons gained or lost, it equals the charge.
Examples
(1)nonmetals
(2) metals
POLYATOMIC IONS AND COMPOUNDS
The octet rule for the formation of compounds states that atoms tend to gain, lose, or share electrons so that the outermost energy level holds or shares four pairs of electrons (an octet – eight electrons).
This tendency leads to an ns2, np6 or noble gas configuration, a particularly stable electronic state (demonstrated by the relative scarcity of noble gas compounds).
Hydrogen is satisfied with 2 electrons in its outermost orbit (duet rule).
IONIC COMPOUNDS
COVALENT COMPOUNDS
Rules (when several atoms are present):
(1)Determine the number of valence electrons present.
(2)Determine the number of electrons needed to satisfy the octet rule with no electron sharing.
(3)Subtract the number of valence electrons from the number of electrons required to satisfy the octet rule with no sharing to find the number of bonding electrons.
(4)Arrange the atoms as symmetrically as possible. Remember that H can bond to only one other atom, as it is capable of sharing only 2 electrons.
(5)Place the number of electrons to be shared between atoms.
(6)Add the remainder of the electrons to complete the octets or duets of all other atoms.
Example
Examples
POLYATOMIC IONS
When drawing e-dot diagrams for polyatomic ions, you do not split the polyatomic ions into their elements.