Download chemical periodicity

CHEMICAL PERIODICITY
ANTOINE LAVOISIER (1790s)
“The father of modern chemistry
The first to recognize the true elements
JOHANN DOBEREINER (1817)
Found that several groups of three elements have similar properties and called it TRIADS
Second element has an atomic mass halfway between 1st and 3rd elements
STANISLAO CANNIZZARO (1860)
Presented an accurate method for measuring relative masses of atoms at the First
International Congress of Chemists in Karlsruhe, Germany
This initiated a search for the relationships between atomic mass and properties of elements
Clearly distinguished between atoms and molecules
JOHN NEWLANDS (1863-1864)
Arranged the elements in order of their increasing atomic masses
Elements were arranged in sequential order, assigning an ordinal number to each
Noted that there is repetition of similar properties every eighth element
Placed seven elements in each group and called it OCTAVES
LAW OF OCTAVES: _______________________________________________
Did not work for all elements
DMITRI MENDELEEV (1869)
Listed 60+ elements in several vertical columns in order of their increasing atomic mass
Noticed a regular recurrence of their physical and chemical properties
____________________ in the table because there were no known elements with the
appropriate properties at that time
1871 - Predicted the physical and chemical properties of 3 missing elements
_________________________
German chemist
Working on element organization at the same time as Mendeleev
Published his work a year later than Mendeleev
WILLIAM RAMSEY 1868-1900
Discovery of noble gases
Proposed a new group in the Periodic Table to accommodate the noble gases discovered
Noble gases were hard to detect due to __________________________________
HENRY MOSELEY (1911-1913)
Mendeleev had a few elements that did not fit
Determined the nuclear charge and atomic number of elements by analyzing the spectra of
38 metals
Rearranged the elements in the table in the order of their __________________________
HARRY D. HUBBARD (1924)
Modernized Mendeleev's periodic table
His first work was published in 1924 known as the "Periodic Chart of the Atoms"
POshikiri / Chemistry 1 of 9 GLENN SEABORG (1940)
Discovered 10 Transuranium elements
Reorganized the table in 1944 by placing the lanthanide/actinide series at
________________ of the table
MODERN PERIODIC TABLE
PERIODIC TABLE
Arrangement of elements in order of their increasing atomic numbers so that
elements with similar properties fall in the same column or group
PERIODIC LAW
The physical and chemical properties of the elements are periodic functions of their
______________________
PERIOD or SERIES
Horizontal rows of the periodic table
7 periods
Denote # of _________________________
GROUP or FAMILY
Vertical columns of the periodic table
Denote # of _________________________
POshikiri / Chemistry 2 of 9 POshikiri / Chemistry 3 of 9 ELECTRON CONFIGURATION & PT
MAIN GROUP ELEMENTS
s- and p-block elements
1A to 8A (1,2,,13,14,15,16,17,18)
REPRESENTATIVE ELEMENTS
Group A elements in the periodic table
1A, 2A, 3A, 4A, 5A, 6A, 7A
s and p sublevels are partially filled
s-Block: Outermost electrons are in the s-sublevel
Alkali Metals – very active metals
Alkaline Earth Metals – also active
Hydrogen and Helium
p-Block
Groups 13 – 18
Electrons add to a p-sublevel only after the s- sublevel in the same energy level is
filled
Boron, carbon, nitrogen, oxygen, halogens,noble gas families
HALOGENS – salt formers, very reactive NM
METALLOIDS – brittle solids with some properties of both metals and nonmetals
NOBLE GASES – inert due to filled s and p orbitals
d-Block
Group B elements in the periodic table/Groups 3-12
Transition Elements – d-block elements with typical metallic properties
Form colored ions
f-Block
Lanthanides: 4f sublevel is being filled
Actinides: 5f sublevel is being filled
POshikiri / Chemistry 4 of 9 PERIODIC PROPERTIES
Shielding Effect
Atomic Radius
Ionization Energy
Electron Affinity
Ionic Radii
Electronegativity
ATOMIC FORCES
NUCLEAR PULL
Theoretical force equal to Z
Z = # of protons
ELECTRON REPULSION
SHIELDING
Inner shells of e- “shield” the outer shell e- from the nucleus’ pull
EFFECTIVE NUCLEAR CHARGE
EFFECTIVE NUCLEAR CHARGE (Zeff) is the
“_______________________________________________.
Outer electrons feel less than full strength of nucleus
because of electrons between them
Z (shielding effect) = Zeff
Example:
Na (Z = 11)
Filled 1s2s2p effectively blocks out 10p+
Valence e- only feels 1 p+
Zeff = 1
Within a period, as Zeff ______________, radius
________________
SHIELDING EFFECT
Decrease in the attraction between outer electrons and the nucleus due to the presence of
other electrons between them
PERIOD TREND
Constant within a period because the _________________________________
GROUP TREND
Increases from top to bottom because of the increase in the main energy levels
ATOMIC RADIUS
_____________________ between the nuclei of identical
atoms that are bonded together
GROUP TREND
Atomic radii of the main group elements increase
down a group
This is due to _________________ in the number of
main energy levels
Because electrons are added further from the
nucleus, there is _____________________.
Size of the atom is generally determined by the number of energy levels
POshikiri / Chemistry 5 of 9 PERIOD TREND
Decreasing trend across a period
Across a period, there is an increase of protons in the nucleus, and also an increase of
electrons in the same main energy level
Size ____________________________________________. Each added electron feels a
greater and greater + charge.
The increase in nuclear charge increases the attraction of the nucleus for the
outermost electrons, pulling it tightly towards the nucleus, causing the size of the
atom to decrease
IONIZATION ENERGY
Energy required to _______________________________
Or how tightly an e- is held by atom
ION – atom or group of bonded atoms that has a positive or negative charge
IONIZATION – process that results in the formation of ion
First ionization energy (IE1)
Minimum amount of energy needed to remove the most loosely held e- from an atom
Ex:
Ca + 590 kJ → Ca1+ + eSecond ionization energy (IE2)
Amount of energy needed to remove the second eEx:
Ca1+ + 1145 kJ → Ca2+ + eIE2 is always greater than IE1
____________________________ to remove e- from +ion than a neutral atom
FACTORS AFFECTING IONIZATION ENERGY
NUCLEAR CHARGE
The larger the nuclear charge, the greater the ionization energy
SHIELDING EFFECT
The greater the shielding effect, the less the ionization energy
RADIUS
The greater the distance between the nucleus and outer electrons, the less the
ionization energy
SUBLEVEL
An electron from a full or half-full sublevel requires additional energy to be removed
GROUP TREND
___________________ down a group
Size of the atom increases as we go down
The outermost electron is farther away
from the nucleus, thus there is decreased
attraction between nucleus and outermost
electrons
Outermost electrons are easily removed and
therefore have a lower ionization energy
PERIOD TREND
Ionization energies of main group elements ___________________________
Nuclear charge increases within a period, which strongly attracts the added electrons
in the same energy level
The atoms become smaller; outermost electrons are held tightly to the nucleus,
making it harder to remove
Slight decrease between IIA and IIIA due to p sublevels having a higher energy so
electrons are easier to remove
POshikiri / Chemistry 6 of 9 Slight decrease between VA and VIA due to paired electrons in the p sublevel. Paired
electrons have greater repulsive forces so they are easier to remove
Noble gases have highest IE (don’t want to let go of electrons)
IE helps predict whether an element is likely to form an
___________________________________.
Metals tend to have ______ IE and form cations, while nonmetals tend to have _____ IE and
form anions – thus forming ionic bonds
Elements with intermediate IE form molecular/covalent compounds by sharing electrons
ELECTRON AFFINITY
Energy change that occurs when a gas phase atom gains an electron to form a gas phase
anion.
A(g) + e- ---> A-(g) E.A. = ∆E
Or atom’s attraction for additional electrons
PERIOD TREND
_____________ as we move left to right across a period
The atoms become smaller and the nuclear charge increases; outermost electrons are
held tightly to the nucleus
Increasingly negative left to right across a period
GROUP TREND
Electron affinity ________________ as we move down a group because of the
increasing atomic size, which decreases attraction between the nucleus and the
outermost electrons
Increasingly positive down a group so electrons add with greater difficulty
Metals have low electron affinities
Nonmetals have high electron affinities
Exceptions with IIA, VA, and VIIIA
Why?
Increased stability
IIA – full s
VA – ½ full p
VIIIA – full s and p
Decreased tendency to gain/lose eIONIC RADII
REMEMBER THAT ALL ATOMS:
Want to have a noble gas configuration
Achieve maximum stability with lowest energy configuration
Gain/lose/share e- to make it happen
Formation of Ions – electrons are removed from the outermost main energy level
CATION (POSITIVE ION)
Smaller than neutral atom because of the loss of outer shell electrons
Na: 1s22s22p63s1
Na+: 1s22s22p6 – same as [Ne]
When a cation is formed from an atom of a transition metal, electrons are always removed
first from the ns orbital and then from the (n – 1)d orbitals.
Fe: [Ar]4s23d6
Mn: [Ar]4s23d5
2+
0
6
6
Fe : [Ar]4s 3d or [Ar]3d
Mn2+: [Ar]4s03d5 or [Ar]3d5
Fe3+: [Ar]4s03d5 or [Ar]3d5
POshikiri / Chemistry 7 of 9 ANION (NEGATIVE ION)
Always larger than neutral atoms because nuclear attraction is spread over additional
electrons
Increased electron-electron repulsion causes the species to become larger
F-: 1s22s22p6 or [Ne]
PERIOD TREND
______________ in the size of positive
ions from left to right across a period
From Group 15, negative ions (which
are much larger in size), gradually
decreases from left to right
For d-block and f-block elements,
electrons are removed from s-sublevel
before d- or f- electrons, resulting in
formation of different kinds of ions for
same element
GROUP TREND
___________ in ionic radii for both
anions and cations as you go down each group
Outer electrons in both cations and anions are in higher energy levels, so there is a
gradual increase of ionic radii
ELECTRONEGATIVITY
Measure of the ability of an atom in a chemical compound to attract electrons
Electronegativity values help predict the type of bonding that can exist between atoms in
compounds
PERIOD TRENDS
___________________ from left to right due to
increase in nuclear pull
Metallic elements - very low EN
Fr & Cs most reactive metals
Nonmetallic elements - have high EN
F - highest EN = most reactive NM
GROUP TRENDS
_________________ from top to bottom within a
group due to increased shielding effect
VALENCE ELECTRONS
Electrons in the _______________________________, which are available to be lost, gained
or shared in the formation of chemical compounds
OXIDATION NUMBERS
The ___________________________________ in a molecule if electrons were transferred
completely in the direction indicated by the _____________________________
Metals are found on the left and center of the periodic table
Their atoms tend to lose electrons and thus have positive oxidation numbers
Transition metals can form as many as 4 different cations because of their complex
electron arrangement
Nonmetals are on the right side of the periodic table
Their atoms tend to gain electrons and thus have ___________________________
POshikiri / Chemistry 8 of 9 OXIDES
Metallic oxides form
________________________
Nonmetallic oxides form
_____________________
______________ – a substance that
has the properties of an acid and a
base
METALLIC CHARACTER
POshikiri / Chemistry 9 of 9