chapter 3 chem

7/28/2019 Chapter 3 Chem

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Jutiz, France Louie R. CN 23

Development of the Periodic Table

The quest for a systematic arrangement of the elements started with the discovery ofindividual elements. By 1860 about 60 elements were known and a method was needed

for organization. In fact many scientists made significant contributions that eventuallyenabled Mendeleev to construct his table. The periodic table did not end withMendeleev but continued to take shape for the next 75 years.

Dobereiners Law of TriadsThe development of the periodic table begins with German chemist Johann Dobereiner(1780-1849) who grouped elements based on similarities. Calcium (atomic weight 40),strontium (atomic weight 88), and barium (atomic weight 137) possess similar chemicalprepares. Dobereiner noticed the atomic weight of strontium fell midway between theweights of calcium and barium:

Ca Sr Ba (40 + 137) 2 = 8840 88 137

Was this merely a coincidence or did some pattern to the arrangement of the elements

exist? Dobereiner noticed the same pattern for the alkali metal triad (Li/Na/K) and thehalogen triad (Cl/Br/I).

Li Na K Cl Br I7 23 39 35 80 127

In 1829 Dobereiner proposed the Law of Triads: Middle element in the triad had atomicweight that was the average of the other two members. Soon other scientists foundchemical relationships extended beyond triads. Fluorine was added to Cl/Br/I group;sulfur, oxygen, selenium and tellurium were grouped into a family; nitrogen,phosphorus, arsenic, antimony, and bismuth were classified as another group.

First Periodic TableIt was a 19th century geologist who first recognized periodicity in the physical propertiesof the elements. Alexandre Beguyer de Chancourtois (1820-1886), professor of geology

at the School of Mines in Paris, published in 1862 a list of all the known elements. Thelist was constructed as a helical graph wrapped around a cylinder--elements with similarproperties occupied positions on the same vertical line of cylinder (the list also includedsome ions and compounds). Using geological terms and published without thediagram, de Chancourtois ideas were completely ignored until the work of Mendeleev.

Law of OctavesEnglish chemist John Newlands (1837-1898), having arranged the 62 known elementsin order of increasing atomic weights, noted that after interval of eight elements similarphysical/chemical properties reappeared. Newlands was the first to formulate theconcept of periodicity in the properties of the chemical elements. In 1863 he wrote apaper proposing the Law of Octaves: Elements exhibit similar behavior to the eighthelement following it in the table.

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Mendeleev's Periodic TableThen in 1869, Russian chemist Dimitri Mendeleev (1834-1907) proposed arrangingelements by atomic weights and properties. Mendeleev's periodic table of 1869contained 17 columns with two partial periods of seven elements each (Li-F & Na-Cl)followed by two nearly complete periods (K-Br & Rb-I).

In 1871 Mendeleev revised the 17-group table with eight columns (the eighth groupconsisted of transition elements). This table exhibited similarities not only in small unitssuch as the triads, but showed similarities in an entire network of vertical, horizontal,and diagonal relationships. The table contained gaps but Mendeleev predicted the ofreceiving the Nobel Prize in chemistry.

Noble GasesLord Rayleigh (1842-1919) and William Ramsey (1852-1916) greatly enhanced theperiodic table by discovering the "inert gases." In 1895 Rayleigh reported the discoveryof a new gaseous element named argon. This element was chemically inert and did notfit any of the known periodic groups. Ramsey followed by discovering the remainder of

the inert gases and positioning them in the periodic table. So by 1900, the periodic tablewas taking shape with elements were arranged by atomic weight. For example, 16goxygen reacts with 40g calcium, 88g strontium, or 137g barium. If oxygen used as thereference, then Ca/Sr/Ba assigned atomic weights of 40, 88, and 137 respectively.

Rayleigh (physics) and Ramsey (chemistry) were awarded Nobel prizes in 1904. Thefirst inert gas compound was made in 1962 (xenon tetrafluoride) and numerouscompounds have followed--today the group is more appropriately called the noblegases.

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Moseley's Periodic LawSoon after Rutherford's landmark experiment of discovering the proton in 1911, HenryMoseley (1887-1915) subjected known elements to x-rays. He was able to derive therelationship between x-ray frequency and number of protons. When Moseley arrangedthe elements according to increasing atomic numbers and not atomic masses, some of

the inconsistencies associated with Mendeleev's table were eliminated. The modernperiodic table is based on Moseley's Periodic Law (atomic numbers). At age 28,Moseley was killed in action during World War I and as a direct result Britain adoptedthe policy of exempting scientists from fighting in wars.

Modern Periodic TableThe last major change to the periodic table resulted from Glenn Seaborg's work in themiddle of the 20th century. Starting with plutonium in 1940, Seaborg discoveredtransuranium elements 94 to 102 and reconfigured the periodic table by placing thelanthanide/actinide series at the bottom of the table. In 1951 Seaborg was awarded theNobel Prize in chemistry and element 106 was later named seaborgium (Sg) in hishonor.

Source: http://web.fccj.org/~ethall/period/period.htm

Dmitri Mendeleev's predicted elements

Professor Dmitri Mendeleev published the first Periodic Table of the Atomic Elements in1869 based on properties which appeared with some regularity as he laid out theelements from lightest to heaviest. When Mendeleev proposed his periodic table, henoted gaps in the table, and predicted that as-yet unknown elements existed withproperties appropriate to fill those gaps.

Prefixes

To give provisional names to his predicted elements, Mendeleev used the prefixes eka-,dvi-, and tri-, from the Sanskrit words for one, two, and three, depending upon whetherthe predicted element was one, two, or three places down from the known element inhis table with similar chemical properties. For example, germanium was calledekasilicon until its discovery in 1886, and rhenium was called dvi-manganese before itsdiscovery in 1926.

Sometimes the eka- prefix is used to refer to some transuranic elements, for exampleeka-lead for ununquadium, eka-radon for ununoctium and eka-caesium.

Nowadays, the prefix eka- (and, more rarely, dvi-) is sometimes used in discussionsabout undiscovered elements, such as untriennium, also known as eka-actinium or dvi-lanthanum.

Current official IUPAC practice is to use a systematic element name based on theatomic number of the element as the provisional name, instead of being based on itsposition in the periodic table as these prefixes require.

Original predictions from 1870The four predicted elements lighter than the rare earth elements, ekaboron (Eb),ekaaluminium (Ea), ekamanganese (Em), and ekasilicon (Es), proved to be goodpredictors of the properties of scandium, gallium, technetium and germanium

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respectively, which each fill the spot in the periodic table assigned by Mendeleev. Initialversions of the periodic table did not give the rare earth elements the treatment nowgiven them, helping to explain both why Mendeleevs predictions for heavier unknownelements did not fare as well as those for the lightest predictions and why they are notas well known or documented.

Ekaboron and scandiumScandium oxide was isolated in late 1879 by Lars Fredrick Nilson; Per Teodor Cleve

recognized the correspondence and notified Mendeleev late in that year. Mendeleevhad predicted an atomic mass of 44 for ekaboron in 1871 while scandium has an atomicmass of 44.955910.

Ekaaluminium and galliumIn 1871 Mendeleev predicted the existence of a yet-undiscovered element he namedeka-aluminium (because of its proximity to aluminium in the periodic table).

Ekamanganese and technetiumTechnetium was isolated by Carlo Perrier and Emilio Segr in 1937, well afterMendeleevs lifetime, from samples of molybdenum that had been bombarded with

deuterium nuclei in a cyclotron by Ernest Lawrence. Mendeleev had predicted anatomic mass of 100 for ekamanganese in 1871 and the most stable isotope oftechnetium is 98Tc.

Ekasilicon and germaniumGermanium was isolated in 1886, and provided the best confirmation of the theory up tothat time, due to its contrasting more clearly with its neighboring elements than the twopreviously confirmed predictions of Mendeleev do with theirs.

Source: http://en.wikipedia.org/wiki/Dmitri_Mendeleev's_predicted_elements

Periodic LawIn chemistry, the periodic law stating that many of the physical and chemical propertiesof the elements tend to recur in a systematic manner with increasing atomic number.Progressing from the lightest to the heaviest atoms, certain properties of the elementsapproximate those of precursors at regular intervals of 2, 8, 18, and 32. For example,the 2d element (helium) is similar in its chemical behavior to the 10th (neon), as well asto the 18th (argon), the 36th (krypton), the 54th (xenon), and the 86th (radon). Thechemical family called the halogens, composed of elements 9 (fluorine), 17 (chlorine),35 (bromine), 53 (iodine), and 85 (astatine), is an extremely reactive family.

Source:

http://dwb.unl.edu/teacher/nsf/c04/c04links/www.fwkc.com/encyclopedia/low/articles/p/p019000875f.html

The parts of the Modern Periodic Table

1. Metals, non-metals, and metalloids:

The periodic table tells you where the metallic, non-metallic, and semi metallicmetals are. To the right of the periodic table, starting to the left of boron (element #5,

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B) you should see a line that looks like a staircase. Elements far to the left of this lineare metals, elements to the far right of this line are non-metals, and elements rightaround the line on either side are semimetals, or metalloids.

To review: Metals are conductors of heat and electricity, malleable, ductile, andgenerally solid. Non-metals may be solids, liquids, or gases, and are poor conductors ofheat and electricity. When they are solids, they are brittle, non-lustrous materials.

Metalloids are solids at standard conditions, and are semiconductors of electricity,making them handy for use in the electronics field. Metalloids have properties betweenthat of metals and non-metals, causing them to have the nickname of "semimetals."

2. The families of the periodic table:

The periodic table consists of a whole bunch of different families which share similarproperties. Families are columns in the periodic table, also referred to as groups. Alkali metals are group 1. They are highly reactive elements with low melting andboiling points. They are light, soft metals. They tend to form ions with a +1 charge. Alkaline earth metals are group 2. They are also reactive, but less so than the alkalimetals. They are light, soft metals, but stronger and denser than the alkali metals. Theytend to form ions with a charge of +2. Transition metals are in groups 3-12. They are less reactive than the alkali andalkaline earth metals, but vary greatly among themselves in reactivity. Generally, theseelements form cations, but the amount of positive charge these elements have dependson what the metals are reacting with. Lanthanides are the metals in the 4f part of the periodic table. They are generallyreactive, shiny metals with various industrial purposes. Like the transition metals, theyform cations with varying amounts of charge. Actinides are metals in the 5f part of the periodic table. Most are radioactive andman-made. Uses of these elements are primarily in the generation of nuclear power orin nuclear explosives. Small amounts of elements such as americium are used in smokedetectors. Chalcogens are group 16 in the periodic table. Starting with oxygen, these elementsare mostly nonmetallic and somewhat electronegative, forming ions with a -2 charge. Halogens are group 17 in the periodic table. These elements are highly reactiveoxidizers, and all form ions with a -1 charge. All are electronegative. All are alsoextremely dangerous, especially when inhaled. Noble gases are group 18 in the periodic table. They basically don't react withanything because they have a stable octet. They used to be called the inert gases, but itwas found a while back that some can form somewhat unstable compounds withhalogens and oxygen. Hydrogen is element #1 in the periodic table. It is unlike any other element, and isfairly reactive. Depending on what it reacts with, it can either form a +1 ion (hydroniumion, or "proton") or a -1 ion (hydride ion) - generally, the hydronium ion is easier to formthan the hydride ion.

3. PeriodIn the periodic table of the elements, elements are arranged in a series of rows (orperiods) so that those with similar properties appear in vertical columns. Elements of thesame period have the same number of electron shells; with each group across a period,the elements have one more proton and electron and become less metallic. Thisarrangement reflects the periodic recurrence of similar properties as the atomic numberincreases. For example, the alkaline metals lie in one group (group 1) and share similarproperties, such as high reactivity and the tendency to lose one electron to arrive at a

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noble-gas electronic configuration. The periodic table of elements has a total of 118elements.

Modern quantum mechanics explains these periodic trends in properties in terms ofelectron shells. As atomic number increases, shells fill with electrons in approximately

the order shown at right. The filling of each shell corresponds to a row in the table.In the s-block and p-block of the periodic table, elements within the same periodgenerally do not exhibit trends and similarities in properties (vertical trends down groupsare more significant). However in the d-block, trends across periods become significant,and in the f-block elements show a high degree of similarity across periods (particularlythe lanthanides).

Periods

1. Chemical elements in the first period

The first period contains fewer elements than any other, with only two, hydrogen andhelium. They therefore do not follow the octet rule. Chemically, helium behaves as a

noble gas, and thus is taken to be part of the group 18 elements. However, in termsof its nuclear structure it belongs to the s block, and is therefore sometimesclassified as a group 2 element, or simultaneously both 2 and 18. Hydrogen readilyloses and gains an electron, and so behaves chemically as both a group 1 and agroup 17 element.

2. Chemical elements in the second period

Period 2 elements involve the 2s and 2p orbitals. They include the biologically mostessential elements besides hydrogen: carbon, nitrogen, and oxygen.3. Chemical elements in the third period

All period three elements occur in nature and have at least one stable isotope. Allbut the noble gas argon are all essential to basic geology and biology.

4. Chemical elements in the fourth periodPeriod 4 includes the biologically essential elements potassium and calcium, and isthe first period in the d-block with the lighter transition metals. These include iron,the heaviest element forged in main-sequence stars and a principal component ofthe earth, as well as other important metals such as cobalt, nickel, copper, and zinc.

Almost all have biological roles.5. Chemical elements in the fifth periodPeriod 5 contains the heaviest few elements that have biological roles, molybdenumand iodine. (Tungsten, a period 6 element, is the only heavier element that has abiological role.) It includes technetium, the lightest exclusively radioactive element.6. Chemical elements in the sixth period

Period 6 is the first period to include the f-block, with the lanthanides (also known asthe rare earth elements), and includes the heaviest stable elements. Many of theseheavy metals are toxic and some are radioactive, but platinum and gold are largelyinert.7. Chemical elements in the seventh period

All elements of period 7 are radioactive. This period contains the heaviest elementwhich occurs naturally on earth, uranium. All of the subsequent elements in the

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period have been synthesized artificially. Whilst some of these (e.g. plutonium) arenow available in tonne quantities, most are extremely rare, having only beenprepared in microgram amounts or less. Some of the later elements have only everbeen identified in laboratories in quantities of a few atoms at a time.

Although the rarity of many of these elements means that experimental results are

not very extensive, periodic and group trends in behaviour appear to be less welldefined for period 7 than for other periods. Whilst francium and radium do showtypical properties of Groups 1 and 2 respectively, the actinides display a muchgreater variety of behaviour and oxidation states than the lanthanides. Initial studiessuggest Group 14 element ununquadium appears to be a noble gas instead of apoor metal, and group 18 element ununoctium probably is not a noble gas. Thesepeculiarities of period 7 may be due to a variety of factors, including a large degreeof spin-orbit coupling and relativistic effects, ultimately caused by the very highpositive electrical charge from their massive atomic nuclei.

Sources:

http://wiki.answers.com/Q/What_are_the_parts_of_the_modern_periodic_tablehttp://en.wikipedia.org/wiki/Period_(periodic_table)

Different Types of Elements in the Periodic Table

Elements that have similar chemical and physical properties end up in the same

column in the periodic table.

Each vertical column of elements in the periodic table is called a group, or family.

The elements in any group of the periodic table have similar physical and chemical

properties. Each group is identified by a number and the letter A or B.

This gives rise to the periodic law: When the elements are arranged in order ofincreasing atomic number, there is a periodic repetition of their physical andchemical properties.

The Group 1A elements are called the alkali metals, and the Group 2A elements are

called the alkaline earth metals.

Most of the remaining elements that are not Group A elements are also metals.

These include the transition metals and inner transition metals, which together make

up the Group B elements.

Copper, gold, and silver are familiar transition metals.

The inner transition metals, which appear below the main body of the periodic table,

are also called the rare-earth elements.

Approximately 80% of all the elements are metals.

The nonmetals occupy the upper-right corner of the periodic table.

Nonmetals are elements that are generally nonlustrous and that are generally poor

conductors of electricity.

Two groups on nonmetals are given special names.

The nonmetals of Group 7A are called the halogens, which include chlorine and

bromine.

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The nonmetals of Group 0 (or 18 or 8A) are known as the noble gases, which are

sometimes called the inert gases because they undergo few chemical reactions.

On your periodic table you will notice a heavy stair-step line.

This line divides the metals from the nonmetals. Most of the elements that border

this line are metalloids, elements with properties that are intermediate betweenthose of metals and nonmetals.

Silicon and germanium are two important metalloids that are used in the

manufacture of computer chips and solar cells.

Classification of the Elements

The periodic table is probably the most important tool in chemistry. Among otherthings, it is very useful for understanding and predicting the properties of theelements.

For example, if you know the physical and chemical properties of one element in agroup or family, you can predict the physical and chemical properties of the otherelements in the same group - and perhaps even the properties of the elements inneighboring groups.

Classifying Elements by Electron Configuration

Of the three major subatomic particles, the electron plays the most significant role

in determining the physical and chemical properties of an element.

The arrangement of elements in the periodic table depends on these properties.

Thus there should be some relationship between the electron configurations of the

elements and their placement in the table.

Elements can be classified into four categories according to their electron

configurations.

1. The noble gases. These are elements in which the outermost s and psublevels are filled.

The noble gases belong to Group 0. The elements in this group aresometimes called the inert gases because they do not participate in manychemical reactions. The electron configurations for the first four noble-gaselements are listed below. Notice that these elements have filled outermost sandp sublevels.

2. The representative elements. In these elements, the outermost s and psublevel is only partially filled.

The representative elements are usually called the Group A elements.

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For any representative element, the group number equals the number ofelectrons in the outermost energy level.For example, the elements in Group 1A (lithium, sodium, etc.) have oneelectron in the outermost energy level.

3. The transition metals. These are metallic elements in which the outermost ssublevel and nearby d sublevel contains electrons.

The transition elements, called Group B elements, are characterized byaddition of electron to the dorbitals.

4. The inner transition metals. These are metallic elements in which theoutermost s sublevel and nearby f sublevel generally contain electrons.The inner transition metals are characterized by the filling offorbitals.

Source: http://www.slideshare.net/tufdaawg/periodic-table-chapter-14

Oxidation number

In coordination chemistry, the oxidation number of a central atom in a coordinationcompound is the charge that it would have if all the ligands were removed along withthe electron pairs that were shared with the central atom. Oxidation numbers are oftenconfused with oxidation states.

The oxidation number is used in the nomenclature of inorganic compounds. It isrepresented by a Roman numeral. The oxidation number is placed either as a rightsuperscript to the element symbol, for example Fe III, or in parentheses after the name ofthe element, iron (III): in the latter case, there is no space between the element nameand the oxidation number.

Source: http://en.wikipedia.org/wiki/Oxidation_number

Some rules in determining the oxidation number of anatom

The oxidation number is the charge an atom in a substance would have if the pairs ofelectrons in each bond belonged to the more electronegative atom. Now this means thatin a compound made up of monatomic ions, like NaCl, in which the bonding pairs dobelong to the more electronegative atom, the oxidation number equals the ionic charge.The sodium ion has an oxidation number of +1 while the chlorine ion has an oxidationnumber of -1. Here are some rules that are helpful in finding the oxidation number of aspecific atom:

The oxidation number of an atom in an elementary substance is 0. This means thatthe oxidation number of an O atom in O2 is 0. The oxidation number of a Group IA atom in any compound is +1; the oxidation

number of a Group IIA atom in any compound is +2. The oxidation number of fluorine is -1 in all of its compounds. The oxidation number of chlorine, bromine, and iodine is -1 in any compound

containing only two elements. The usual oxidation number of oxygen in a compound is -2. The major exceptions of

this rule are peroxides, like H2O2, which have an oxidation number of -1.

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http://en.wikipedia.org/wiki/Atomhttp://en.wikipedia.org/wiki/Coordination_compoundhttp://en.wikipedia.org/wiki/Coordination_compoundhttp://en.wikipedia.org/wiki/Ligandhttp://en.wikipedia.org/wiki/Electron_pairhttp://en.wikipedia.org/wiki/Oxidation_statehttp://en.wikipedia.org/wiki/Inorganic_nomenclaturehttp://en.wikipedia.org/wiki/Roman_numeralhttp://en.wikipedia.org/wiki/Atomhttp://en.wikipedia.org/wiki/Coordination_compoundhttp://en.wikipedia.org/wiki/Coordination_compoundhttp://en.wikipedia.org/wiki/Ligandhttp://en.wikipedia.org/wiki/Electron_pairhttp://en.wikipedia.org/wiki/Oxidation_statehttp://en.wikipedia.org/wiki/Inorganic_nomenclaturehttp://en.wikipedia.org/wiki/Roman_numeral


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The oxidation number of hydrogen in most compounds is +1. The sum of the oxidation numbers in a compound is always zero. For something that

is an ion consisting of two atoms (a polyatomic ion), the oxidation numbers add up tothe charge on the ion.

Source: http://library.thinkquest.org/19957/redox/oxidationnumbody.html

Valence Electrons

In chemistry, valence electrons are the electrons of an atom that can participate in the

formation of chemical bonds with other atoms. Valence electrons are the "own"

electrons, present in the free neutral atom, that combine with valence electrons of other

atoms to form chemical bonds. In a single covalent bond both atoms contribute one

valence electron to form a shared pair. For main group elements, only the outermost

electrons are valence electrons. In transition metals, some inner-shell electrons are also

valence electrons.

Valence electrons are important in determining how the atom reacts chemically with

other atoms. Atoms with a complete (closed) shell of valence electrons (corresponding

to an electron configuration s2p6) tend to be chemically inert. Atoms with one or two

valence electrons more than a closed shell are highly reactive because the extra

electrons are easily removed to form positive ions. Atoms with one or two valence

electrons fewer than a closed shell are also highly reactive because of a tendency either

to gain the missing electrons and form negative ions, or to share electrons and form

covalent bonds.

Valence electrons have the ability, like electrons in inner shells, to absorb or release

energy in the form of photons. This gain or loss of energy can trigger an electron to

move (jump) to another shell or even break free from the atom and its valence shell.

When an electron absorbs energy in the form of one or more photons, then it moves to

a more outer shell depending on the amount of energy gained. (See also : electrons in

an excited state). When an electron loses energy (photons), then it moves to a more

inner shell.

The number of valence electrons

The number of valence electrons of an element is determined by its periodic table group

(vertical column) in which the element is categorized. With the exception of groups 312

(transition metals), the number within the unit's place identifies how many valence

electrons are contained within the elements listed under that particular column.

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* The general method for counting valence electrons is generally not useful for transition

metals. Instead the modified d electron count method is used.

** Except forhelium, which has only two valence electrons.

Periodic table groupValence

electrons

Group 1 (I) (alkali metals) 1

Group 2 (II) (alkaline earth

metals)2

Groups 3-12 (transition metals) See note *

Group 13 (III) (boron group) 3

Group 14 (IV) (carbon group) 4

Group 15 (V) (nitrogen group) 5

Group 16 (VI) (chalcogens) 6

Group 17 (VII) (halogens) 7

Group 18 (noble gases) 8**

Source: http://en.wikipedia.org/wiki/Valence_electron

Periodic Variations

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Observe your Periodic Table carefully across a period and within a group and you will

observe this.

From Top to Bottom From Left to Right

Atomic Number Increases Increases

Atomic Mass Increases Increases

Atomic Radius

Atomic Radius is one-half the distance between two nuclei in two adjacentidentical atoms and is measured in terms ofpicometer(pm)

Atomic radius could either be metallic radius orcovalent radius. Metallic radiusis a term used for the atomic radius of metallic elements while covalent radius isa term used for the atomic radius of nonmetallic elements.

The atomic radius can be identified by the strength of attraction of the nucleus ofan atom and the outermost or valence electrons.

The strength of attraction is called nuclear charge. The larger the value of the nuclear charge of an atom, the smaller is its atomic

radius. Larger nuclear charge indicates that the hold of the nucleus on the outermost or

valence electron is strong. The larger the value of the valence electron, the larger is the nuclear charge of

the atom and the smaller is its atomic radius. Atomic radius increases from top to bottom or down a group of the Periodic Table

and decreases from left to right or across a period of the Periodic Table.

Ionic Radius

Ionic radius is the radius of a cation or an anion in an ionic compound.An atom tends to change in size as it becomes a cation or anion. Metals give up electrons to form a positively charged ion (cation). Nonmetals gain electrons to form a negatively charged ion (anion).An atom gives up or gains electrons to be stable.As metal forms a cation, its radius decreases because of a lesser electron-

electron repulsion which is a result of the removal of electron. Therefore, a cationhas a smaller ionic radius compared with anion.

As nonmetal forms an anion, its radius increases because of a greater electron-electron repulsion which is a result of the addition of electrons. Therefore, an

anion has a bigger ionic radius compared with a cation.

Ionization Energy

Ionization energy is the minimum energy required to remove an electron from anisolated atom or ion in its ground state and is usually measured in terms ofkilojoules per mole (kJ/mol).

One electron is removed from the atom one at a time.

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When one electron in an atom is removed, the energy needed is called firstionization energy.

Metals have low ionization energies compared with nonmetals that is why metalsreadily give up electrons forming a cation.

Ionization energy increases from left to right or across a period of the PeriodicTable and decreases from top to bottom or down a group of the Periodic Table.

Electron Affinity

Electron Affinity is the energy change when an atom in its ground state gains anelectron forming an anion and is usually measured in terms of kilojoules per mole(kJ/mol).

Nonmetals have high electron affinity that is why nonmetals always form ananion.

Electron affinity increases from left to right or across a period of the Periodictable and from bottom to top or up a group of the Periodic Table.

Electronegativity

Electronegativity is the ability of an atom in a chemical bond to attract electronstoward it and is usually measured in terms of electronvolt (eV).

Electronegativity is related to ionization energy and electron affinity. High electron affinity would mean a greater ability to pick up electrons easily

while high ionization energy would mean a lesser possibility to lose an electron. If an atom has high ionization energy and high electron affinity therefore it also

has high electronegativity. Electronegativity increases from left to right or across a period of the Periodic

Table and it decreases from top to bottom or down a group of the PeriodicTable.

Metallic Character

Metallic character is a chemical property associated with elements classed asmetals.

These are elements that have the tendency to give up electrons and formpositive ion.

In the Periodic Table, metallic character increases down any group and across aperiod from right to left.

Nonmetallic Character

Nonmetallic character is a chemical property associated with elements classedas nonmetals.

These are elements that have the tendency to gain electrons and form negative

ion. In the Periodic Table, nonmetallic character increases up any group and across a

period from left to right.

Source: http://www.scribd.com/doc/21540877/Trends-and-Periodic-Variation-of-

Elements-in-the-Periodic-Table

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Blocks of Elements

A block of the periodic table of elements is a set of adjacent groups. The term appearsto have been first used (in French) by Charles Janet.[1] The respective highest-energyelectrons in each element in a block belong to the same atomic orbital type. Each block

is named after its characteristic orbital; thus, the blocks are: s-block p-block d-block f-block g-block (hypothetical)

The block names (s, p, d, f. and g) are derived from the quality of the spectroscopiclines of the associated atomic orbitals: sharp, principal, diffuse and fundamental, therest being named in alphabetical order. Blocks are sometimes called families.

The following is the order for filling the "subshell" orbitals, according to the Aufbauprinciple, which also gives the linear order of the "blocks" (as atomic number increases)

in the periodic table:

1s, 2s, 2p, 3s, 3p, 4s, 3d, 4p, 5s, 4d, 5p, 6s, 4f, 5d, 6p, 7s, 5f, 6d, 7p, ...

For discussion of the nature of why the energies of the blocks naturally appear inthis order in complex atoms, see atomic orbital and electron configuration.

The "periodic" nature of the filling of orbitals, as well as emergence ofthe s, p, d and f "blocks" is more obvious, if this order of filling is given in matrixform, with increasing principal quantum numbers starting the new rows ("periods") inthe matrix. Then, each subshell (composed of the first two quantum numbers) isrepeated as many times as required for each pair of electrons it may contain. Theresult is a compressed periodic table, with each entry representing two successive

elements:

1s

2s 2p 2p 2p

3s 3p 3p 3p

4s 3d 3d 3d 3d 3d 4p 4p 4p

5s 4d 4d 4d 4d 4d 5p 5p 5p

6s (4f) 5d 5d 5d 5d 5d 6p 6p 6p

7s (5f) 6d 6d 6d 6d 6d 7p 7p 7p

Source: http://en.wikipedia.org/wiki/Block_(periodic_table)

Predicting Properties

There are 88 naturally occurring elements. Chemistry would be a nearly impossiblesubject to study if there were not some relationships between the elements. Chemists

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http://en.wikipedia.org/wiki/Periodic_tablehttp://en.wikipedia.org/wiki/Periodic_table_grouphttp://en.wikipedia.org/wiki/Charles_Janethttp://en.wikipedia.org/wiki/Block_(periodic_table)#cite_note-0http://en.wikipedia.org/wiki/Atomic_orbitalhttp://en.wikipedia.org/wiki/S-blockhttp://en.wikipedia.org/wiki/P-blockhttp://en.wikipedia.org/wiki/D-blockhttp://en.wikipedia.org/wiki/F-blockhttp://en.wikipedia.org/wiki/Extension_of_the_periodic_table_beyond_the_seventh_periodhttp://en.wikipedia.org/wiki/Atomic_orbitalhttp://en.wikipedia.org/wiki/Aufbau_principlehttp://en.wikipedia.org/wiki/Aufbau_principlehttp://en.wikipedia.org/wiki/Atomic_orbitalhttp://en.wikipedia.org/wiki/Electron_configurationhttp://en.wikipedia.org/wiki/Periodic_tablehttp://en.wikipedia.org/wiki/Periodic_table_grouphttp://en.wikipedia.org/wiki/Charles_Janethttp://en.wikipedia.org/wiki/Block_(periodic_table)#cite_note-0http://en.wikipedia.org/wiki/Atomic_orbitalhttp://en.wikipedia.org/wiki/S-blockhttp://en.wikipedia.org/wiki/P-blockhttp://en.wikipedia.org/wiki/D-blockhttp://en.wikipedia.org/wiki/F-blockhttp://en.wikipedia.org/wiki/Extension_of_the_periodic_table_beyond_the_seventh_periodhttp://en.wikipedia.org/wiki/Atomic_orbitalhttp://en.wikipedia.org/wiki/Aufbau_principlehttp://en.wikipedia.org/wiki/Aufbau_principlehttp://en.wikipedia.org/wiki/Atomic_orbitalhttp://en.wikipedia.org/wiki/Electron_configuration


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several centuries ago noticed that the properties of elements seem to reoccurperiodically. At first it was though that this periodicity was a function of atomic weight butmodern scientists have shown that it varies based on atomic number (#of protons) . Asa result they formulated the Periodic Law.

The properties of an element are a periodic (repeating) function of their atomic

numbers.

Mendeleev was one of the first to use such a law (although he thought the propertiesvaried due to atomic weight) to create a table which organized the elements. He placedelements that had similar properties underneath each other in columns. These columnsare known as chemical families.

We understand today that the properties vary according to atomic number and thereason for the similar chemical and physical properties is due to the electron structure ofthe atom. Therefore elements in the same column in the periodic table have both similarproperties and electron arrangement.

One of the chemical properties that are similar is the formula of the compounds created

with other elements.

Source:

http://www.saskschools.ca/~bisstchem/modules/module2/lesson2/predictingproperties.h

tm

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chapter 3 chem

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