For Grade X

By:

Widy Alif Putra Said

4301412095

Semarang State University

2015

Chemistry Atom and Periodic Table of Element

Basic Competency :

3.2 Analyzing the development of atomic model.

3.3 Analyzing the atomic particles based on Bohr atomic theory and Mechanical Quantum

theory.

3.4 Analyzing the relationship electron configuration and orbital diagram to determine the

position of element in periodic table and element periodic prroperties.

4.2 Processing and analyzing the development of atomic model.

4.3 Processing and analyzing the atomic particles based on Bohr atomic theory and

Mechanical Quantum theory.

4.4 Providing the result of analyzing the relationship electron configuration and orbital

diagram to determine the position of element in periodic table and element periodic

properties.

Indicators :

1. Comparing the development of element periodic table to identify its majority and minority.

2. Explaining basic of grouping elements.

3. Determining basic particles (proton, electron and netron).

4. Determining electron configuration and valence electron.

5. Determining atomic relative mass based on periodic table.

6. Grouping the element into isotop, isobar and isoton.

7. Grouping the element into metal, non metal and metaloid.

8. Analyzing table, graphik to determine regularity of atomic radius, ionization energy, electron

affinity and electronegativity.

Atomic Theory Development

1. Democritus

460 BC - Greek philosopher proposes the existence of the atom

He pounded materials until he made them into smaller and smaller parts

He called them atoma which is Greek for “indivisible”.

His Theory:

All atoms:

Are small hard particles

Are made of a single material formed into different shapes and sizes

Are always moving, and they form different materials by joining together

2. John Dalton

Dalton's atomic theory contains five basic assumptions:

All matter consists of tiny particles called atoms. Dalton and others imagined the atoms

that composed all matter as tiny, solid spheres in various stages of motion.

Atoms are indestructible and unchangeable. Atoms of an element cannot be created,

destroyed, divided into smaller pieces, or transformed into atoms of another element.

Dalton based this hypothesis on the law of conservation of mass as stated by Antoine

Lavoisier and others around 1785.

Elements are characterized by the weight of their atoms. Dalton suggested that all atoms

of the same element have identical weights. Therefore, every single atom of an element

such as oxygen is identical to every other oxygen atom. However, atoms of different

elements, such as oxygen and mercury, are different from each other.

In chemical reactions, atoms combine in small, whole-number ratios. Experiments that

Dalton and others performed indicated that chemical reactions proceed according to atom

to atom ratios which were precise and well-defined.

When elements react, their atoms may combine in more than one whole-number ratio.

Dalton used this assumption to explain why the ratios of two elements in various

compounds, such as oxygen and nitrogen in nitrogen oxides, differed by multiples of

each other.

John Dalton's atomic theory was generally accepted because it explained the laws of

conservation of mass, definite proportions, multiple proportions, and other observations.

Although exceptions to Dalton's theory are now known, his theory has endured reasonably well,

with modifications, throughout the years.

3. J J. Thomson

1897 - English chemist and physicist; discovered 1st subatomic particles

His Theory:

o Atoms contain negatively charged particles called electrons and positively

charged matter.

o Created a model to describe the atom as a sphere filled with positive matter with

negative particles mixed in

Referred to it as the plum pudding model

Plum Pudding Model or Raisin Bun

ModelProposed by J.J.

Thomson

4. Ernest Rutherford

1912 - New Zealand physicist discovered the nucleus

His Theory:

Small, dense, positively charged particle present in nucleus called a proton

Electrons travel around the nucleus, but their exact places cannot be described.

5. Niels Bohr Atomic Models

Bohr atomic model suggests that the atom consists of a

nucleus is very small and is surrounded by a positively

charged electron has a negative charge that orbits

This is the picture theory Bohr model of the atom.

Nuclear ModelProposed by Ernest

Rutherford

Figure 1. Bohr Atom Model

Niels Bohr Bohr's atomic theory put forward in 1915. Due to the Bohr model of the atom

is a modification (development) of the Rutherford atomic model, some chemists call the

Rutherford-Bohr theory of the atom. Although Bohr's atomic theory had been developed, but in

fact the Bohr model of the atom still has weaknesses. However, some points of the Bohr model

of the atom can be accepted. Unlike Dalton's atomic theory and atomic theory of Rutherford,

Bohr atomic theory excellence can explain the Rydberg constant for hydrogen line emission

spectra. That's one of the advantages Niels Bohr atomic theory.

Bohr atomic model shaped like a solar system, with electrons which are in circulation

path (orbit) around the positively charged nucleus is very small in size.

The force of gravity in the solar system can be illustrated mathematically as the Coulomb

force between a nucleus (nuclei) are positively charged with negatively charged electrons.

The sound of the Bohr atomic theory postulates

Bohr atomic theory can be explained as follows:

1. Electrons circling the atomic nucleus in specific orbits were circular. These orbits are

often referred to as electron shells are represented by the notation K, L, M, N ... etc

that are berututan accordance with n = 1, 2, 3, 4 ... and so on.

2. Electrons in each orbit has a higher specific energy with the increasing size of the

growing circular orbit or the value of n. This energy is quantized and prices permitted

by the stated price quantized angular momentum of electrons n (h/2π) with n = 1, 2, 3,

4 ... and so on.

3. During the orbit, the electron does not radiate energy and said in a stationary state.

The existence of a stationary electron in orbit is maintained by the force of

electrostatic attraction of electrons by atomic nuclei is balanced by the centrifugal

force of the motion of electrons.

4. Electrons can move from one orbit to another orbit which has a higher energy when

the electrons absorb energy in accordance with the magnitude of the energy

difference between the two orbits is concerned, and vice versa when the electrons

move into an orbit that has a lower energy would emit radiation energy line spectra

are observed as the magnitude corresponding to the energy difference between the

two orbits is concerned.

5. Atoms in a molecule is said in the state of the ground level (ground state) when the

electrons occupy orbits so as to give the lowest total energy. And when the electrons

occupy orbits which gives a higher energy than the energy levels of atoms in

essentially said the excited level (excited state). Atoms in the ground state is more

stable than in the excited state.

Model Hidrogen Bohr

The simplest example of the Bohr model of the hydrogen atom (Z = 1) or a hydrogen-like

ions (Z> 1), which has negatively charged electrons surround the positively charged nucleus.

Electromagnetic energy is absorbed or released when an electron moves from one track to

another track. The radius of the track increases as n2, where n is the principal quantum number.

Transition from 3 to 2 produces the first line of the Balmer series. For hydrogen (Z = 1) will

produce a photon with a wavelength of 656 nm (red light).

Weakness Bohr Atomic Theory

Although considered to be a revolutionary, but still found weaknesses Bohr atomic theory

are:

Breaking the Heisenberg uncertainty principle because the electron has a radius and a

known trajectory.

Bohr atomic model has a value of the angular momentum of the ground state trajectory is

wrong.

Weak explanation of atomic spectra predicted greater.

Unable to predict the relative intensities of the spectral lines.

Bohr atomic model can not explain the structure of the spectral lines are well.

Unable to explain the Zeeman effect.

6. Quantum Mechanics Model of the Atom

An explanation of a more complete atomic structure necessary to know more details

about the structure of electrons in the atom. Complete atomic model should be able to explain the

mystery of the Zeeman effect and is suitable for many electron atoms. Two of these symptoms

can not be explained by the Bohr model of the atom.

Effect Zeeman

Atomic line spectrum is observed when an electric current is passed through gas in a tube

of gas prod. Additional lines observed in the emission spectrum when excited atoms placed in an

external magnetic field. One line in the spectrum of visible emission lines as three lines (with

two additional lines) in the spectra of atoms when placed in a magnetic field. Breaking up a line

into multiple lines in a magnetic field is known as the Zeeman effect.

Figure 2. Separation of atomic spectral lines in a magnetic field

Zeeman effect can not be explained using the Bohr model of the atom. As such, required

a more complete model of the atom and the more general to explain the Zeeman effect and the

spectrum of many electron atoms.

Model Atomic model of Quantum Mechanics

Previously we have discussed about the wave-particle duality which states that an object

can behave either as a wave or a particle. the atomic scale, the electrons can be reviewed as a

wave phenomenon that does not have a particular position in space. The position of an electron is

represented by a probability or chance of finding the electron in the largest space.

To get a complete and general description of the structure of the atom, the wave-particle

duality principle used. Here the electron motion is described as a wave phenomenon. Newton's

equation of dynamics that was originally used to describe the motion of electrons is replaced by

the Schrodinger equation for the wave function of the electron states. Atomic model which is

based on this principle is called the atomic model of quantum mechanics.

Figure 3. Atomic model of quantum mechanics

To further investigate the structure, we need to know a term called the "quantum number".

There are four types of quantum numbers, is;

1. First Quantum Numbers (n)

2. Quantum Numbers Azimuth (l)

3. Magnetic Quantum Numbers (m)

4. The spin quantum number (s)

First Quantum Numbers (n)

Is a number that states the energy level or shell electrons are:

Shells number Shells Total electron max

(2n)2

(n=1) K 2(1)2 = 2

(n=2) L 2(2)2 = 8

(n=3) M  2(3)2 = 18

(n=4) N   2(4)2 = 32

Quantum Numbers Azimuth (l)

     Numbers stating the location of a subshell

Shell number Shell Value subshell Subshell

(n=1) K 0 S

(n=2) L 0, 1 P

(n=3) M 0, 1 ,2 D

(n=4) N 0, 1, 2, 3 F

subshell Number of Orbital

S 1

P 3

D 5

F 7

Magnetic Quantum Numbers (m)

Numbers stating it was the location of an orbital

Subshell value (m)

S 0

P -1,0,+1

D -2,-1, 0, +1, +2

F -3, -2,-1, 0 +1, +2, +3

Spin Quantum Numbers

Spin quantum number is needed to explain the anomalous Zeeman effect. This anomaly in the

form of splitting of spectral lines into more lines than expected. If the Zeeman effect is caused by

the presence of an external magnetic field, the anomalous Zeeman effect caused by the rotation

of the electron on its axis. Rotation or spin of electrons produces an intrinsic angular momentum

of electrons. Spin angular momentum also has two different orientations, namely spin up and

spin down. Each electron spin orientations have different energies thin so it looks as separate

spectral lines of atoms in the magnetic field derived from the electron spin.

Electron spin quantum number is represented by a separate so-called magnetic spin quantum

number (or so-called spin only). Value of the spin quantum number should only be one of two

values + ½ or - ½. ifmsadalah spin quantum number, angular momentum components of the z-

axis direction is written as

Spin up is expressed with

Spin down is expressed with

Periodic Table of Element Development

Chemists have always looked for ways of arranging the elements to reflect the

similarities between their properties. The modern periodic table lists the elements in order of

increasing atomic number (the number of protons in the nucleus of an atom). Historically,

however, relative atomic masses were used by scientists trying to organise the elements. This was

mainly because the idea of atoms being made up of smaller sub-atomic particles (protons,

neutrons and electrons) had not been developed. Nevertheless, the basis of the modern periodic

table was well established and even used to predict the properties of undiscovered elements long

before the concept of the atomic number was developed.

Table formation

Ask most chemists who discovered the periodic table and you will almost certainly get the answer Dmitri Mendeleev. Certainly Mendeleev was the first to publish a version of the table that we would recognise today, but does he deserve all the credit?

A number of other chemists before Mendeleev were investigating patterns in the properties of the elements that were known at the time. The earliest attempt to classify the elements was in 1789, when Antoine Lavoisier grouped the elements based on their properties into gases, non-metals, metals and earths. Several other attempts were made to group elements together over the coming decades. In 1829, Johann Döbereiner recognised triads of elements with chemically similar properties, such as lithium, sodium and potassium, and showed that the properties of the middle element could be predicted from the properties of the other two

It was not until a more accurate list of the atomic mass of the elements became available at a

conference in Karlsruhe, Germany in 1860 that real progress was made towards the discovery of the modern periodic table.

This area of the website celebrates the work of many famous scientists whose quest to learn more about the world we live in and the atoms that make up the things around us led to the periodic table as we know it today.

Alexandre-Emile Béguyer de Chancourtois

Alexandre Béguyer de Chancourtois. Reproduced courtesy of Annales des mines, Paris.

Can France claim the first periodic table?  Probably not, but a French Geology Professor

made a significant advance towards it, even though at the time few people were aware of it. 

Alexandre Béguyer de Chancourtois was a geologist, but this was at a time when

scientists specialised much less than they do today. His principal contribution to chemistry was

the 'vis tellurique' (telluric screw), a three-dimensional arrangement of the elements constituting

an early form of the periodic classification, published in 1862. 

The telluric screw plotted the atomic weights of the elements on the outside of a cylinder,

so that one complete turn corresponded to an atomic weight increase of 16. As the diagram

shows, this arrangement means that certain elements with similar properties appear in a vertical

line.  Although the telluric screw did not correctly display all the trends that were known at the

time, de Chancourtois was the first to use a periodic arrangement of all of the known elements,

showing that similar elements appear at periodic atom weights.

The vis tellurique from De

Chancourtois’s original

publication (right) and a

copy drawn out with

modern symbols (left).

John Newlands

John Newlands. Reproduced courtesy of the Library and Information Centre, Royal Society of

Chemistry.

John Newlands was British; his father was a Scottish Presbyterian minister.  He was

educated by his father at home, and then studied for a year (1856) at the Royal College of

Chemistry, which is now part of Imperial College London.  Later he worked at an agricultural

college trying to find patterns of behaviour in organic chemistry.  However, he is remembered

for his search for a pattern in inorganic chemistry.

Just four years before Mendeleev announced his periodic table, Newlands noticed that

there were similarities between elements with atomic weights that differed by seven.  He called

this The Law of Octaves, drawing a comparison with the octaves of music.  The noble gases

(Helium, Neon, Argon etc.) were not discovered until much later, which explains why there was

a periodicity of 7 and not 8 in Newlands table.  Newlands did not leave any gaps for

undiscovered elements in his table, and sometimes had to cram two elements into one box in

order to keep the pattern.  Because of this, the Chemical Society refused to publish his paper,

with one Professor Foster saying he might have equally well listed the elements alphabetically.

Even when Mendeleev had published his table, and Newlands claimed to have discovered

it first, the Chemical Society would not back him up.  In 1884 he was asked to give a lecture of

the Periodic Law by the Society, which went some way towards making amends.  Finally, in

1998 the Royal Society of Chemistry oversaw the placing a blue commemorative plaque on the

wall of his birthplace, recognising his discovery at last.

The blue commemorative plaque placed at

Newlands’ birthplace, declaring him the

“discoverer of the Periodic Law for the

chemical elements”.

Julius Lothar Meyer

Julius Lothar Meyer . Reproduced courtesy of the Library and Information Centre, Royal Society

of Chemistry.

Meyer trained at Heidelberg University under Bunsen and Kirchhoff, as did Mendeleev.

So the two scientists would certainly have known each other although neither was aware of all

the work done by the other. Meyer's roots, however, were firmly in Germany.  Meyer was just

four years older than Mendeleev, and produced several Periodic Tables between 1864-1870.

His first table contained just 28 elements, organised by their valency (how many other

atoms they can combine with). These elements were almost entirely main group elements, but in

1868 he incorporated the transition metals in a much more developed table.  This 1868 table

listed the elements in order of atomic weight, with elements with the same valency arranged in

vertical lines, strikingly similar to Mendeleev’s table.  Unfortunately for Meyer, his work wasn’t

published until 1870, a year after Mendeleev’s periodic table had been published. Even after

1870, Meyer and Mendeleev were still unaware of each other’s work, although Meyer later

admitted that Mendeleev had published his version first.

Meyer did contribute to the development of the periodic table in another way though.  He

was the first person to recognise the periodic trends in the properties of elements, and the graph

shows the pattern he saw in the atomic volume of an element plotted against its atomic weight.

A modern version of Meyer’s

graph demonstrating the periodic

trends in the atomic volume of the

elements, plotted against atomic

weight.

Dmitri Mendeleev

Dmitri Mendeleev. Reproduced courtesy of the Library and Information Centre, Royal Society of

Chemistry.

As we have seen, Mendeleev was not the first to attempt to find order within the

elements, but it is his attempt that was so successful that it now forms the basis of the modern

periodic table.

Mendeleev did not have the easiest of starts in life. He was born at Tobolsk in 1834, the

youngest child of a large Siberian family. His father died while he was young, and so his mother

moved the family 1500 km to St. Petersburg, where she managed to get Dmitri into a “good

school“, recognising his potential. In his adult life he was a brilliant scientist, rising quickly in

academic circles. He wrote a textbook, Chemical Principles, because he couldn’t find an

adequate Russian book.

Mendeleev discovered the periodic table (or Periodic System, as he called it) while

attempting to organise the elements in February of 1869. He did so by writing the properties of

the elements on pieces of card and arranging and rearranging them until he realised that, by

putting them in order of increasing atomic weight, certain types of element regularly occurred.

For example, a reactive non-metal was directly followed by a very reactive light metal and then a

less reactive light metal. Initially, the table had similar elements in horizontal rows, but he soon

changed them to fit in vertical columns, as we see today.

Not only did Mendeleev arrange the elements in the correct way, but if an element

appeared to be in the wrong place due to its atomic weight, he moved it to where it fitted with the

pattern he had discovered. For example, iodine and tellurium should be the other way around,

based on atomic weights, but Mendeleev saw that iodine was very similar to the rest of the

halogens (fluorine, chlorine, bromine), and tellurium similar to the group 6 elements (oxygen,

sulphur, selenium), so he swapped them over.

The real genius of Mendeleev’s achievement was to leave gaps for undiscovered

elements. He even predicted the properties of five of these elements and their compounds. And

over the next 15 years, three of these elements were discovered and Mendeleev’s predictions

shown to be incredibly accurate. The table below shows the example of Gallium, which

Mendeleev called eka-aluminium, because it was the element after aluminium. Scandium and

Germanium were the other two elements discovered by 1886, and helped to cement the

reputation of Mendeleev’s periodic table.

The final triumph of Mendeleev’s work was slightly unexpected. The discovery of the

noble gases during the 1890s by William Ramsay initially seemed to contradict Mendeleev’s

work, until he realised that actually they were further proof of his system, fitting in as the final

group on his table. This gave the table the periodicity of 8 which we know, rather than 7 as it had

previously been. Mendeleev never received a Nobel Prize for his work, but element 101 was

named Mendelevium after him, an even rarer distinction.

  Eka-aluminium (Ea) Gallium (Ga)

Atomic weight About 68 69.72

Density of solid 6.0 g/cm³ 5.9 g/cm³

Melting point Low 29.78°C

Valency 3 3

Method of discovery Probably from its spectrum Spectroscopically

Oxide

Formula Ea2O3, density 5.5

g/cm3.  Soluble in both acids

and alkalis

Formula Ga2O3, density 5.88

g/cm3.  Soluble in both acids and

alkalis

A comparison of Mendeleev’s predicted “Eka-aluminium” and Gallium, discovered by Paul

Emile Lecoq in 1875

A commemorative stamp showing

Mendeleev and some of his original notes

about the Periodic Table

Henry Moseley

Henry Moseley. Reproduced courtesy of the Library and Information Centre, The Royal Society

of Chemistry.

It wasn’t until 1913, six years after Mendeleev’s death that the final piece of the puzzle

fell into place. The periodic table was arranged by atomic mass, and this nearly always gives the

same order as the atomic number.  However, there were some exceptions (like iodine and

tellurium, see above), which didn’t work. Mendeleev had seen that they needed to be swapped

around, but it was Moseley that finally determined why.

He fired the newly-developed X-ray gun at samples of the elements, and measured the

wavelength of X-rays given.  He used this to calculate the frequency and found that when the

square root of this frequency was plotted against atomic number, the graph showed a perfect

straight line. He’d found a way to actually measure atomic number.  When the First World War

broke out, Moseley turned down a position as a professor at Oxford and became an officer in the

Royal Engineers.  He was killed by a sniper in Turkey in August 15, and many people think that

Britain lost a future Nobel prize winner.

Within 10 years of his work, the structure of the atom had been determined through the

work of many prominent scientists of the day, and this explained further why Moseley’s X-rays

corresponded so well with atomic number.  The idea behind the explanation is that when an

electron falls from a higher energy level to a lower one, the energy is released as electromagnetic

waves, in this case X-rays.  The amount of energy that is given out depends on how strongly the

electrons are attracted to the nucleus.  The more protons an atom has in its nucleus, the more

strongly the electrons will be attracted and the more energy will be given out.  As we know,

atomic number is also known as proton number, and it is the amount of protons that determine

the energy of the X-rays.

After years of searching, at last we had a periodic table that really worked, and the fact

that we still use it today is testament to the huge achievement of these and many other great

minds of the last two centuries of scientific discovery.

Periodic Table of Element

Period - A horizontal row in the periodic table.

The energy levels of the s and p orbitals are numbered by the row in which they are located.

e.g.  The 2s orbital is in the second row (Li and Be) and the 3p orbitals are in the third row (Al,

Si, P, S, Cl, Ar)

The d orbitals are placed one row below their energy level.

e.g.  The 3d orbitals are in the fourth row

Group - A vertical column, or family, in the periodic table.

Numbered in two ways:

1 2 3 4 5 6 7 8 9 10 11 12 13 14 15 16 17 18

IAII

A

III

B

IV

B

V

B

VI

B

VII

BVIIIB IB

II

B

III

A

IV

A

V

A

VI

A

VII

A

VIII

A

H He

Li Be B C N O F Ne

Na Mg Al Si P S Cl Ar

K Ca Sc Ti V Cr Mn Fe Co Ni Cu Zn Ga Ge As Se Br Kr

Rb Sr Y Zr Nb Mo Tc Ru Rh Pd Ag Cd In Sn Sb Te I Xe

Cs Ba La Hf Ta W Re Os Ir Pt Au Hg Tl Pb Bi Po At Rn

Fr Ra Ac Unq Unp Unh Uns Uno Une

Group 1 (IA) - Alkali Metals (excluding H)

1. Li (lithium), Na (sodium), K (potassium), Rb (rubidium), Cs (cesium), and Fr (francium)

All form hydroxides (e.g.  NaOH)

2. All are active metals

3. Activity increases as you move down the column

React violently when they come into contact with water

4. All have one valence electron

  Li: [He] 2s1   Rb: [Kr] 5s1

Na: [Ne] 3s1 Cs: [Xe] 6s1

K: [Ar] 4s1 Fr: [Rn] 7s1

5. All lose one electron to form cations with a charge of +1 e.g.  Li+, Na+, and K+

Group 2 (IIA) - Alkaline Earth Metals

1. Be (beryllium), Mg (magnesium), Ca (calcium), Sr (strontium), Ba (barium), and Ra

(radium)

All but Be and Mg are active metals

2. Activity increases as you move down the column

Ca, Sr, and Ba react violently when they come into contact with water

3. All have two valence electrons

  Be: [He] 2s2   Sr: [Kr] 5s2

Mg: [Ne] 3s2 Ba: [Xe] 6s2

Ca: [Ar] 4s2 Ra: [Rn] 7s2

4. Tend to lose two electrons to form cations with a charge of +2 e.g.  Be2+, Mg2+, and Ca2+

Alkaline earth metals are less active than their adjacent alkali metals e.g.  Be is less active

than Li; Mg is less active than Na

Groups 3-12 (IIIB - VIIIB, IB & IIB) - Transition Metals

1. All have valence electrons in the d orbitals

Form compounds in which they have various oxidation numbers

2. Ag (silver) generally forms Ag+ (+1 oxidation state)

Zn (zinc) generally forms Zn2+ (+2 oxidation state)

3. Brass is an alloy of copper and zinc

Group 13 (IIIA)

1. B (boron) - only element in the group that is not a metal; has semimetal and nonmetal

characteristics.

Al (aluminum) - fairly active metal, third most abundant in the earth's crust.

2. Loses three electrons to form Al3+

3. Forms compounds in which it has an oxidation state of +3

4. Other metals - Ga (gallium), In (indium), and Tl (thallium) - very scarce active metals

Group 14 (IVA)

1. C (carbon) - nonmetal

2. Elemental forms of carbon include:

Graphite (crystalline) - Strong bonds between atoms within planes resulting in extremely

high melting and boiling points.  Weaker bonds connecting the planes which account for

the soft texture of graphite.

Diamond (crystalline) - Hardest naturally occurring substance with extremely high

melting and boiling points.  Atoms arranged in a tetrahedral array with strong C-C bonds.

Charcoal - Results from heating wood without oxygen present

Coke (amorphous) - More structured than other amorphous forms of carbon; made from

coal.

Carbon Black (amorphous) - Formed by burning natural gas or other carbon compounds

in a limited amount of air

3. Has strong C-C single bonds, C=C double bonds, and C C triple bonds.

Forms covalent bonds with other elements

Can form double and triple bonds with other nonmetals

4. S (silicon) and Ge (germanium) - semimetals

5. Silicon is the second most abundant element in the earth's crust

6. Sn (tin) and Pb (lead) - less reactive metals

7. Both form compounds in which their oxidation states are +2 or +4

Bronze is an alloy of tin and copper

Group 15 (VA)

1. N (nitrogen) and P (phosphorus) - nonmetals

2. Nitrogen is found in its elemental form at room temperature as a diatomic gas (N2).

Nitrogen makes up approximately 80% of earth's atmosphere by volume.

In compounds, the oxidation of nitrogen can range from -3 to +5.

Haber process - mixing N2 and H2 gases at 200 to 300 atm and 400 C to 600 C over a

finely divided iron catalyst to produce NH3

Pure elemental phosphorus is white phosphorus (P4).  It is highly reactive and combusts

with air at room temperature but is unreactive with water.

Red phosphorus is formed when white phosphorus is heated and is much less reactive

than white phosphorus.

Phosphorus can expand its valence (outermost) shell to hold more than eight electrons

(can store extra electrons in the 3d orbitals).

N N triple bonds are much stronger than P P.

P-P single bonds are stronger than N-N single bonds.

3. As (arsenic) and Sb (antimony) - semimetals

Bi (bismuth) - metal

Group 16 (VIA)

1. (oxygen), S (sulfur), and Se (selenium) - nonmetals

2. Oxygen is the most abundant element on earth, making up approximately 45% of the

earth's crust (by weight), 85% of the oceans (by weight) and 20% of the atmosphere (by

volume).

Oxygen is generally diatomic (O2) in its elemental form, but ozone (O3) is an allotrope of

O2.

3. At concentrations above 1 ppm, ozone is toxic.

Ozone can absorb ultraviolet (UV) radiation from the sun, and serves as a filter in the

atmosphere.

4. Oxygen is a very strong oxidizing agent, weaker only to fluorine.

Oxygen generally takes on a -2 oxidation number in compounds.

5. In peroxide (O22-), oxygen has a -1 oxidation number, e.g. H2O2, hydrogen peroxide.

6. Elemental sulfur is a yellow solid at room temperature with a cyclical molecular structure

(S8).

O=O double bonds are much stronger than S=S double bonds.

S-S single bonds are stronger than O-O bonds.

Sulfur can expand its valence (outermost) shell, to hold more than eight electrons (can

store extra electrons in the 3d orbitals).

In compounds, sulfur can have oxidation numbers ranging from -2 to +6.

The prefix thio- is given to compounds in which an S atom replaces an O atom.

7. Te (tellurium) and Po (polonium) - semimetals

Group 17 (VIIA) - Halogens

1. F (fluorine), Cl (chlorine), Br (bromine), I (iodine), and At (astatine)

All are nonmetals except for At which is a semimetal

All are diatomic in their elemental form

All gain one electron to form anions with a charge of -1.

2. e.g.  F-, Cl-, and Br-

3. In compounds, their oxidation numbers range from -1 to +7.

None are found in nature in their elemental forms; instead they are found as salts of the

halide ions.

Properties (at room temperature)

4. F2 - highly toxic, colorless gas, most reactive element known.

5. So reactive that it can even form compounds with noble gases (once thought to be inert)

Used in manufacturing Teflon, (C2F4)n

Used to make freons which are used in refrigerators

6. Cl2 - highly toxic, pale yellow-green gas, strong oxidizing agent

7. Used commercially as a bleaching agent an as a disinfectant

8. Br2 - reddish-orange liquid with a bad odor

9. Its name, bromine, is derived from the Greek bromos which means "stench"

Used in preparing fire-extinguishing agents, sedatives, insecticides, and antiknock agents

for gasoline

10. I - deep purple solid with a metallic-like luster

11. Sublimes directly into a violet gas (I2) from the solid phase when heated without passing

through the liquid phase

Used as a disinfectant, catalysts, drugs, and dyes

AgI (silver iodide) is used in photography

Iodine deficiency in the human body can lead to a goiter, a swelling of the thyroid gland.

Group 18 (VIIIA) - Noble (Rare) Gases

1. He (helium - "sun"), Ne (neon - "new one"), Ar (argon - "lazy one"), Kr (krypton -

"hidden one"), Xe (xenon - "stranger"), and Rn (radon)

Mistakenly labeled as "inert gases" until about 30 years ago because it was thought that

these gases did not react with anything.

In 1962, Neil Barlett isolated the first compound containing a noble gas: [Xe+][PtF6-]

2. Since then, compounds containing Kr, Xe, and Rn have been isolated, but none

containing He, Ne, or Ar.

3. The oxidation numbers of the rare gases in compounds include +1, +2, +4, +6, and +8.

4. Noble gases have filled valence (outermost) shells.

Electron Configuration

Example : 16S = 1s2 2s2 2p6 3s2 3p4 S2- = 1s2 2s2 2p6 3s2 3p6 11Na+= 1s2 2s2 2p6

How to Determine Group and Period based on Electron Configuration

Example : 16S = 1s2 2s2 2p6 3s2 3p4 Group: the latest number electron in sub shell (VI A)

Periode: the latest number of shell (3)

Periodic Properties:

1. Atomic Radius : amount that when to be summed it will give inter nucleus distace among the

relevance atom

2. Electron Affinity: the energy where atom in the gaseous state is release when it captures an

electron so that became ion -1 charge unit.

3. Ionization energi: the minimum energy where atom in the gaseous state is to release electron.

4. Electronegativity: the ability of an atom to attract electron toward itself in a chemical bond.