CHAPTER 2Atoms molecules and Ions

Dr. Rajani Srinivasan

Tarleton State University

Lecture Presentation

Contents

Atomic weights

Discovery of atomic structure

Modern View of atomic structure

Atomic Theory of Matter

The Periodic Table

Contents

Molecules and Molecular compound

Ions and Ionic Compound

Nomenclature of Inorganic Compounds

Nomenclature of Organic compounds

Atomic theory of Matter

Daltons atomic theory ( John Dalton 1766-1844)

1. Elements are composed of very small particles called

2. All atoms of given element are identical but different from atoms of different elements

ATOMS

Oxygen Chlorine

Daltons atomic theory3.Atoms can neither be created nor be destroyed in Chemical reactions .

Example: Oxygen cannot be changed to nitrogen by chemical Reaction

Law of conservation of Mass Law of Multiple proportion

Daltons atomic theory

4. Compounds are formed when atoms of more than one element combine; a given compound always has the same relative number and kind of atoms.

H H +

O

H HO

Law of constant composition

Laws

• Law of constant composition

In a given compound the relative numbers and kinds of atoms constant • Law of conservation of mass

Total mass of the materials before and after the chemical reaction remains the same

H H +

O

H HO

Mass of hydrogen = 2*1 = 2 + Mass of Oxygen = 16 = 18 amu

Laws

• Law of multiple proportion

“ If the two elements A and B combine to form more than one compound, the masses of B that can combine with a given mass of A is in a ratio of small whole number.”

Example H2O2 and H2O ( thus the ratio of oxygen in the compounds is 2:1)

2:32 2:16

1:16 1: 8

Discovery of Atomic structureSubatomic Particles:• Scientists found by various experiments that

“Atoms are not indivisible”• They are composed of several other particles

smaller than atoms so they were called Subatomic Particles.

• Thus the structure of atoms was discovered through series of landmark experiments.

• Atoms are composed of Electron, Proton and Neutron

Electron

• Discovered by Joseph John Thomson (J.J. Thomson) in 1897

• Electrons were discovered by Cathode Ray tube Experiment.

http://www.youtube.com/watch?v=dehxVQAUqBs

Negatively charged particles

Rays originated from cathode (-ve electrode) and travelled to anode (+ve electrode) so were called Cathode Rays.

Charge to mass ratio

1) These rays were deflected by electric and magnetic fields proving that Cathode rays are –vely Charged particles which was Later named as “ELECTRON”

2) Thus by calculating the strength of the magnetic field Thompson calculated the Charge to mass ratio of Electron = 1.76 108 coulombs/gram (C/g).Coulomb is the SI unit for electric charge

Charge on the ElectronMillikan’s Oil drop experiment was used to measure the charge on the electron(Robert Millikan 1909) • He allowed the drops of oil to fall and

by varying the voltage he discovered that charge on any drop of oil was equal to 1.602 ×10 -19 C

• Thus by using the charge on electron and charge /mass ratio; mass on the electron was calculated using the following formula

Electric mass = 1.602 ×10 -19 C = 1.76 ×10 8 c/g

= 9.10 ×10-28 g

Since electrons had mass so they were considered –vely Charged particles

Radioactivity

• Spontaneous emission of Radiation from a compound/ atom is called Radioactivity.

• Discovered by Henry Becquerel in Uranium compound.

• His students Marie Curie and her husband Pierre Curie began experiments in isolating the components of radiation

• Earnest Rutherford discovered the components of Radioactive radiation

Radioactive Radiation

Three types of radiation were discovered by Ernest Rutherford: particles particles rays

Radioactivity

– particles – have +ve charge, fast moving, attracted towards negative plate, has a charge of +2, has a mass of 7400 times than that of electron

– particles – have –ve charge, fast moving, attracted towards positive plate. Have a charge of -1

– rays – high energy radiation like x-rays, no particles and no charge

Nuclear Model of atom

• Structure of atom as described by J.J. Thomson

• He described the structure as positive sphere of matter with electrons embedded in it.

NucleusRutherford’s Gold foil Experiment:

Important Observations1) Most of the rays passed

undeflected2) Some were scattered in large

angles3) Some scattered in the same

directions from which they had come

Observations 1 and 2 did not agree with Thomson's Model

Rutherford’s Model of Atom

1) In an atom most of the space is Empty

2) Atom has a very small but dense positively charged center called “NUCLEUS” ( indicated by large deflecting angles)

3) Electrons move around the nucleus (indicated by few rays which passed undeflected)

Protons and Neutrons

• Rutherford discovered Protons in 1919

They are positively charged particles with mass of 1.0073 amu• James Chadwick discovered “Neutrons”(1932)

They are neutral particles and has mass equal to that of Proton

http://www.youtube.com/watch?v=HnmEI94URK8t

Subatomic Particles

Modern atomic Structure 1) Atom is made up of Nucleus in the center

2) Nucleus has both Neutrons and Protons

3) Electron revolve round the nucleus

4) An Atom is electrically Neutral ( it has equal number of Electrons and Protons)

6) Diameter of most of the atoms (1-5 ◦A)

Structure of the nucleus

Atomic Number

• The Number of Protons or electrons in an Atom is called “Atomic Number”

• Represented by “Z”

Atomic Mass

• Total of Protons and Neutrons in an atom is called Atomic Mass or Mass Number

• Represented by amu (Atomic Mass Unit)• Atomic Mass Unit – defined by assigning a mass of

exactly 12 amu of 12C isotope of carbon

Isotopes• Atoms of the same element having different masses.• They have same number of Protons and Electrons

but they differ in the number of Neutrons.• Different isotopes of the same element exists in

different abundances

Atomic weight

• Calculation of atomic mass of atoms• Average atomic mass of the element is called

“Atomic Weight” of the element • It is based on the relative abundances of the

isotopes

Atomic weight = ∑ [(Isotope mass) × (fractional isotope abundance)]

of the overall isotope of the element

Atomic weight

• Calculate the atomic weight of Carbon if it is composed of 98.93 % 12 C and 1.07% of 13C

Example :

First convert the % into fractions then use the formula

0.9893* 12 amu+ 0.0107*13 amu = 12amu

Mass spectrometer is used to measure atomic and molecular masses

Periodic Table

North American style of writing

IUPAC Nomenclature

• The periodic table is a systematic catalog of the elements.• Elements are arranged in the increasing order of atomic number.• IUPAC- International Union of Pure and Applied chemistry

Parts of Periodic table

Columns = Groups

Rows = Periods

Elements in the same group have similar properties

Repeating patterns of reactivities could be found in

the period

Important groups

• Nonmetals are on the right side of the periodic table (with the exception of H).

• Metalloids border the stair-step line (with the exception of Al, Po, and At).Properties of the Metalloids fall between metals and non-metals

• Metals are on the left side of the chart.

Molecules and Molecular CompoundMolecules

Most of the element present in the nature are in Molecular form except a few

Like Noble gasses: He, Ar, Kr etc.

Example: Oxygen exists As O2 or O3

Halogens Like Chlorine exists as Cl2 , Br2 etc.

Such type of molecules which are composed of two atoms to form a molecule are called “Diatomic Molecule”

Compounds composed of two or more atoms are called molecular compounds and usually contains Non-metals

Diatomic molecules

• These seven elements occur naturally as molecules containing two atoms:– Hydrogen- H2

– Nitrogen- N2

– Oxygen- O2

– Fluorine- F2

– Chlorine- Cl2

– Bromine- Br2

– Iodine- I2

Triatomic molecule

• Most common example Ozone represented by Chemical Formula O3

• Oxygen and Ozone are made up of same atom but they exhibit very different physical and chemical properties

• O2 is life saving gas ,odorless• O3 is very toxic , has a very pungent smell

Chemical formulas• Molecular formulas give the exact number of

atoms of each element in a compound. • Empirical formulas give the lowest whole-

number ratio of atoms of each element in a compound

Example : Molecular formula for glucose C6H12O6

Empirical formula- Divide the subscripts by largest common factor we get CH2O

Chemical formulas

• Structural formulas show the order in which atoms are bonded.

• Perspective drawings also show the three-dimensional array of atoms in a compound.

• Ball and stick Model This represents the accurate angles in which atoms are attached in a molecule

• Space filling Model Depicts how molecules will look like when they are scaled up

Types of Compounds

1) Ionic compounds

2) Molecular compounds

Ionic compoundsIONS: “Ions” are formed when an atom either gains an electron or loses an electron.

IONS

CATION(+vely Charged

Ion )

ANION(-vely Charged

ION)

How are ions Formed????

CATIONS ANIONS

Example of Na

Z = 11P = 11or 11 +ve ChargeE = 11 or 11 -ve charge

If it loses an electron What will be the

P = E =

Total extra charge on the Na = +1Thus Forming Na + (sodium ion )

1110

Example of Cl

Z= 17P= 17E= 17

If it gains an electron, Then

P= E =

Total No. of Extra Charge on Cl = -1Thus forming Cl- ( Chlorine ion)

1817

IONIC compounds

• When +ve ions and –ve ions combine to form compound they are called “IONIC COMPOUND” The bond between them are called IONIC BOND.

• POLYATOMIC COMPOUNDs: When more than 2 atoms combine to from compounds they are said be Polyatomic ionic Compounds

Example NH3+ and SO4

2- = Na2SO4

IONIC COMPOUNDS

Superscripts

Writing Formulas

Because compounds are electrically neutral, one can determine the formula of a compound this way:

• The charge on the cation becomes the subscript on the anion.

• The charge on the anion becomes the subscript on the cation.

• If these subscripts are not in the lowest whole-number ratio, divide them by the greatest common factor.

Common Cations

Common Anions

Molecular compounds

• When two or more atoms combine together to form compounds.

• The bonding between them is not ionic but

Co-valent. • Usually when compounds are formed between

two non metals it will be a molecular compound

• Example CH4 (Methane) or H2O2

Why are compounds formed

Answer :To complete the Octet

Octet = Compare the Nobel gasses they have 8 electron in their outermost shell so they are self sufficient

“All the other element wants to have 8 electrons in their outermost shell to become self sufficient in that process they gain or lose electron and from compounds”

Nomenclature of Inorganic and Organic compound

NOMENCALARE

Name To Call

ORGANIC INORGANIC

Substances made up of Carbon and Hydrogen in

combination with Oxygen , Nitrogen and Hydrogen

Example – CH4, CH2COOHCH2CONH2

All the rest of the compounds

Example- N2, CO2, NH3

NH4OH

In organic compound Nomenclature• Naming of Cation

1) If it has a single charge. Example Na+

Sodium ion

2) If the Cation has more than one Charge

Fe2+ Iron (II) ion

Fe 3+ Iron(III) ion

Name of the element ion

Name of the element Charge in Roman numeral in the parenthesis

ion

ANIONS

• Elemental Anion : The end of the element gets replaced by ide

Example : Cl-1 = Chloride ( Chlorine replaced by ide)

O2- = Oxide ( Oxygen replaced by ide)

Polyatomic Anion: OH-1 = Hydroxide ion

Oxyanions Polyatomic anions containing Oxygen are called “Oxyanions”The names end in -ate or -ite• When there are two oxyanions involving the same

element:– The one with fewer oxygens ends in -ite.– The one with more oxygens ends in -ate.

• NO2− : nitrite; SO3

2− : sulfite

• NO3− : nitrate; SO4

2− : sulfate

Oxyanions

The one with the second fewest oxygens ends in -ite.– ClO2

− : chloriteThe one with the second most oxygens ends in -ate.

– ClO3− : chlorate

The one with the fewest oxygens has the prefix hypo- and ends in -ite.– ClO− : hypochlorite

The one with the most oxygens has the prefix per- and ends in -ate.– ClO4

− : perchlorate

Oxyanions

• Oxyanions derived from adding H+ are named by adding the name hydrogen or dihydrogen

• CO32-....... HCO3

-…. Hydrogen Carbonate

• SO32--…………….. HSO3

-- Hydrogen Sulphite

ACIDS

Hydrogen containing substance

OR

Substances which yield Hydrogen when dissolved in water• Composed of anions connected to enough H+

ion to neutralize the anions charge• Eg: Cl- + H+ HCl (Hydrochloric acid)

ACIDS• If the anion in the acid

ends in -ide, change the ending to -ic acid and add the prefix hydro- .– HCl: hydrochloric acid– HBr: hydrobromic acid– HI: hydroiodic acid

• If the anion in the acid ends in -ite, change the ending to -ous acid.– HClO: hypochlorous acid

– HClO2: chlorous acid

ACIDS

If the anion in the acid ends in -ate, change the ending to -ic acid.

HClO3: chloric acidHClO4: perchloric acid

Binary molecular CompoundsMolecules with two atoms forming a compound

Example : Cl2O

1) Name of the element nearer to the metals are named first except when Oxygen, Chlorine , Bromine and Iodine are present (except F) Then oxygen is written last.

Cl2O- dichloro monoxide2) If both elements from the same group then lower one is named first Name of the second element ends in –ideClF- Chrloro monoflouride3) Greek prefixes are used to indicate the number of atoms.4) The prefix mono is never used with the first element 5) When the prefix ends with an a or o the second element begins with a vowel N2O5- Di nitrogen tetra oxide

Organic Compounds

• Simplest organic compounds are called ALKANES

• Each Carbon atoms are attached to 4 atoms• They are named based on number of Carbon

and hydrogen atoms .

H-C-H H3C-CH3 H3C-CH2-CH3 C8H18

H

H

Methane Ethane Propane Octane CnH2n+2

Derivatives of Alkanes

• When one or more atoms of Hydrogen is replaced by functional groups like –OH, -COOH are called derivatives of Alkanes

ISOMERS

• Compounds with same molecular formula but different structural formulas are called ISOMERS.

C5H12 can have two different structural arrangement

“The unique property of organic compound is its ability to form Long Chain of Carbon –Carbon bonds.”----- CATENTION