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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
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)
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
• 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
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
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.
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