Physical Chemistry

Every field of science involves taking measurements, understanding them, and communicating them to others. In other words, we all have to speak the same basic language. Whether you are a chemist, a physicist, a biologist, an engineer, or even a medical doctor, you need a consistent way of communicating size, mass, shape, temperature, time, amount, energy, power, and speed.

The International System of Units (abbreviated SI, from the French Système international d’unités) is the metric system used in science, industry, and medicine. Depending on your age and geographic location, you might be very familiar with the “imperial” system, which includes units such as gallons, feet, miles, and pounds. The imperial system is used for “everyday” measurements in a few places, such as the United States. But in most of the world (including Europe) and in all scientific circles, the SI system is in common use.

The SI is made up of 7 base units that define the 22 derived units with special names and symbols.

History of the SI System

The SI units of measurement have an interesting history. Over time they have been refined for clarity and simplicity.

1. The meter (m), or metre, was originally defined as 1/10,000,000 of the distance from the Earth’s equator to the North Pole measured on the circumference through Paris. In modern terms, it is defined as the distance traveled by light in a vacuum over a time interval of 1/299,792,458 of a second.
2. The kilogram (kg) was originally defined as the mass of a liter (i.e., of one thousandth of a cubic meter). It is currently defined as the mass of a platinum-iridium kilogram sample maintained by the Bureau International des Poids et Mesures in Sevres, France.
3. The second (s) was originally based on a “standard day” of 24 hours, with each hour divided in 60 minutes and each minute divided in 60 seconds. However, we now know that a complete rotation of the Earth actually takes 23 hours, 56 minutes, and 4.1 seconds. Therefore, a second is now defined as the duration of 9,192,631,770 periods of the radiation corresponding to the transition between the two hyperfine levels of the ground state of the cesium-133 atom.
4. The ampere (A) is a measure of the amount of electric charge passing a point in an electric circuit per unit time. 6.241×10¹⁸ electrons, or one coulomb, per second constitutes one ampere.
5. The kelvin (K) is the unit of the thermodynamic temperature scale. This scale starts at 0 K. The incremental size of the kelvin is the same as that of the degree on the Celsius (also called centigrade) scale. The kelvin is the fraction 1/273.16 of the thermodynamic temperature of the triple point of water (exactly 0.01 °C, or 32.018 °F).
6. The mole (mol) is a number that relates molecular or atomic mass to a constant number of particles. It is defined as the amount of a substance that contains as many elementary entities as there are atoms in 0.012 kg of carbon-12.
7. The candela (cd) was so named to refer to “candlepower” back in the days when candles were the most common source of illumination (because so many people used candles, their properties were standardized). Now, with the prevalence of incandescent and fluorescent light sources, the candela is defined as the luminous intensity in a given direction of a source that emits monochromatic radiation of frequency 540⋅1012

Derived units – are defined in terms of the seven base quantities via a system of quantity equations. The SI derived units for these derived quantities are obtained from these equations and the seven SI base units

For ease of understanding and convenience, 22 SI derived units have been given special names and symbols,

Ex: volume = L x L x L = m³ or cm³ 

Density=massvolume=kgm3 or gmm3  

Force: unit N (Newton) = kg.m/sec²

Conversion of one unit to another example

                       m³ = m x m x m = (100cm) x (100) cm x (100) cm = 10⁶ cm³

 Density kgm3=1000gm106cm3=10-3 gmcm3

Measurement of data: when conducting experiments we have to measure & report data. There are certain norms for reporting this data.   

All very big and small values are expressed in exponents

PRECISION AND ACCURACY

•  Precision indicates how closely repeated measurements match each other.
•  Accuracy indicates how closely a measurement matches the correct or expected value.
•  A result is valid only if it is both accurate and precise.

EXAMPLE

• If the true value for a result is 8.00 kg and a student  “A” takes two measurements and reports the results as 7.95 kg and 7.93 kg. These values are precise as they are close to each other but are not accurate.
• Another student “B” repeats the experiment and obtains 7.94 kg and 8.05 kg as the results for two measurements. These observations are neither precise nor accurate.
• When a third student “C” repeats these measurements and reports 8.01 kg and 7.99 kg. These values are both precise and accurate.

SIGNIFICANT FIGURES

Significant figures are meaningful digits which are known with certainty. The uncertainty is indicated by writing the certain digits and the last uncertain digit. Thus, if we write a result as 11.2 ml, we say the 11 is certain and 2 is uncertain and the uncertainty would be ±1 in the last digit.

Unless otherwise stated, an uncertainty of ± 1 in the last digit is always understood.

RULES FOR DETERMINING THE NUMBER OF SIGNIFICANT FIGURES

1. All non-zero digits are significant
2. Zeros preceding the first non-zero digit are not significant
3.  Zeros between two non-zero digits are significant.
4.  Zeros at the end or right of the number are significant provided they are on the right side of the decimal point. But, if otherwise, the zeros are not significant
5. Counting numbers of objects, for example, 2 balls or 20 eggs, have infinite significant figures as these are exact numbers

EXAMPLE

ADDITION AND SUBTRACTION OF SIGNIFICANT FIGURES

During addition and subtraction, the result cannot have more digits to the right of the decimal point than either of the original numbers.

MULTIPLICATION AND DIVISION OF SIGNIFICANT FIGURES

In multiplication and division with significant figures, the answer cannot have more significant figures than either of the original numbers.

                                Example               Example

                                  2.12                         1.2

                             +  6.1                    x     1.3

                                 0.012                       1.56

                                 8.232             Ans.   1.6

                      Ans.= 8.2                             

ROUNDING OFF THE SIGNIFICANT FIGURES

1. If the rightmost digit to be removed is more than 5, the preceding number is increased by one. For example, in figure 2.486 if we have to remove 6, we have to round it to 2.49

2. If the rightmost digit to be removed is less than 5, the preceding number is not changed.

For example, in figure 6.664 if 4 is to be removed, then the result is rounded of to 6.66.

3. If the rightmost digit to be removed is 5, then the preceding number is not changed, if it is an even number but it is increased by one if it is an odd number. For example, if 2.35 is to be rounded by removing 5, we have to increase 3 to 4 giving 2.4 as the result. However, if 2.25 is to be rounded off it is rounded off to 2.2

Ex.:  A student performs a titration with different burettes and finds titre values of 25.2 mL, 25.25 mL, and 25.0 ml. The number of significant figures in the average titre value is

(IIT adv. 2010, integer type)

Sol. During addition and subtraction, the result cannot have more digits to the right of the decimal point than either of the original numbers                                                                            Answer:  3

Ex.:  If the value of Avogadro number is 6.023 × 10²³ mol^–1 and the value of Boltzmann constant is 1.380 × 10^–23 J K^–1, then the number of significant digits in the calculated value of the universal gas constant is                                                                                                          (IIT adv. 2014, integer type)

Sol. K = R/N_A  R=KN_A

R= 6.023 × 10²³ × 1.380 × 10^–23 J.mol^–1.k^–1= 8.31174 J.mol^–1.k^–1. There are 4 significant figures in each term. (4) Hence, these be 4 significant figure in R                                                  Answer. 4

INTRODUCTION

The word .atom has been derived from the Greek word atomio which means .un-cutable or non-divisible.

DALTON’S ATOMIC THEORY

The matter is composed of small indivisible particles called atoms (from Greek word atomio, meaning indivisible). In 1808 Dalton proposed the following theory

1. Matter consists of atoms, which cannot be divided further.

2. All the atoms of a given element have identical mass. Atoms of different elements differ in mass.

3. Atoms combine in a fixed ratio to form Compounds.

4 The atoms cannot be created or destroyed.

SUCCESS OF DALTON’S ATOMIC THEORY

1. Atoms are the smallest part of matter
2. Atoms cannot be created of destroyed
3. Atoms of same element are similar in mass
4. Atoms combine with each other in simple ratios.

FAILURES OF DALTON’S ATOMIC THEORY

It could explain lows of chemical combination, as on today we know that all four points are not correct.

CATHODE RAYS

 CATHODE RAYS ARE NEGATIVELY CHARGED

 CATHODE RAYS TRAVEL IN STRAIGHT LINE.

CATHODE RAYS ARE PARTICLES

CHARGE /MASS (E/M) RATIO 

In 1897 JJ. Thomson determined the e/m value of cathode rays.

Thomson proved that e/m ratio is same, whatever material is of plates or gas filled in the tube

Ex.:  Which of the following is never true for cathode rays? 

1. They possess kinetic energy
2. They are electromagnetic waves
3. They produce heat
4. They produce mechanical pressure          Answer (c)

 MILLIKAN’S OIL DROP EXPERIMENT

 MASS OF ELECTRON

From Millikan's oil drop experiment e = 1.6022 x 10^-19 coulombs

   

 

 

ANODE RAYS

 

The smallest and lightest positive ion was obtained from hydrogen and was called proton. This characterized in 1919.

Later the presence of electrically neutral particle was found in the atom. These particles were discovered by Chadwick (1932) by bombarding a thin sheet of beryllium by α-particles, then electrically neutral particles having a mass slightly greater than that of the protons was emitted. He named these particles as neutrons.

Thomson Model of Atom

Plum pudding model OR

Water melon model 

 FUN FACTS JJ THOMSON (NOBEL PRIZE IN 1906)

.STOICHIOMETRY AND VOLUMETRIC

The word stoichiometry is derived from two Greek words - stoicheion (meaning element) and metron (meaning measure).

Stoichiometry, thus, deals with the calculation of masses (sometimes volumes also) of the reactants and the products involved in a chemical reaction.

How many moles of methane are required to produce 22g CO₂ (g) after combustion?

Solution

According to the chemical equation.

CH₄ (g) +2O₂ (g) → CO₂ (g) + 2H₂O (g)

44g CO₂ (g) is obtained from 16 g CH₄ (g).

1 mol CO₂ (g) is obtained from 1 mol of CH₄ (g)]

 

= 0.5 mol CO₂ (g)

Hence, 0.5 mol CO₂ (g) would be obtained from 0.5 mol CH₄ (g) or 0.5 mol of CH₄ (g) would be required to produce 22 g CO₂ (g)

Ex.:- At 300 K and I atmospheric pressure, 10mL of a hydrocarbon required 55mL of O₂ for complete combustion and 40 mL, of CO₂ is formed The formula of the hydrocarbon is        (2019 Main, 10 April)

(a) C₄H₇Cl                                             (b) C₄H₆

(c) C₄H₁₀                                               (d) C₄H₈

Solution: In eudiometry,

  

 

 

 

Given, (i) VCO₂ = 10x = 40mL Þ x = 4

 

 

          

 

 

So, the hydrocarbon (C_x H_y) is C₄H₆         Answer (b)

LIMITING REAGENT

Many a time, the reactions are carried out when the reactants are not present in the amounts as required by a balanced chemical reaction. In such situations, one reactant is in excess over the other. The reactant which is present in the lesser amount gets consumed after sometime and after that no further reaction takes place whatever be the amount of the other reactant present. Hence, the reactant which gets consumed, limits the amount of product formed and is, therefore, called the limiting reagent.

IMPORTANT: Product formed is calculated by limiting reagent

Limiting reagent

2A + 3B → 4C

If we have 3 mole of A of 4 mole of B, find out mole of C farmed

2A + 3B → 4C

   3      4

3/2x  3 = 4.5 > 4 so B is limiting reagent

3 mole of B form 4 mole C

4 mole of B will form

       Answer_: 5.33 moles

Ex.:- If 0.50 mole of BaCl₂ is mixed with 0.20 mole of Na₃PO₄, the maximum number of moles of Ba₃ (PO₄)₂ that can be formed is                                                                                      (1981, 1M)

(a) 0.70                                                 (b) 0.50

(b) 0.20                                                 (d) 0.10

Solution: The balanced chemical reaction is

In this reaction, 3 moles of BaCl₂ combines with 2 moles of Na₃PO₄. Hence, 0.5 mole of BaCl₂ require

 

Since, available Na₃PO₄ (0.2 mole) is less than required mole (0.33), it is the limiting reactant and would determine the amount of product Ba₃(PO₄)₂

∵ 2 moles of Na₃PO₄ gives 1 mole Ba₃ (PO₄)₂

       

 Answer (d)

MOLE CONCEPT

• One mole is the amount of a substance that contains as many particles or entities as there are atoms in exactly 12 g (or 0.012 kg) of the C¹²
• 12gm of C¹² contains 6.022 x10²³ atoms
• The number 6.022 x10²³ is called Avogadro’s constant or Avogadro’s number.
• The mass of one mole of a substance in grams is called its molar mass.
• The molar mass in grams is numerically equal to the atomic/molecular/formula mass in u.(u is the unified mass)

CONVERSION OF MASS & VOLUME TO MOLES

Ex. Which has maximum number of atom?                                                              (2003, 1M)

(a) 24 g of C (12)                                             (b) 56 g of Fe (56)

(c) 27 g of Al (27)                                            (d) 108 g of Ag (108)

Solution: Number of atoms = Number of moles

Numbers of atoms in 24g C= 2412 ×NA=2NA

Number of atoms in 56 g of Fe= 5656NA=NA  

Number of atoms in 27 g of Al= 2727NA=NA  

Number of atoms in 108 of Ag= 108108NA=NA  

Hence, 24g of carbon has the maximum number of atoms.                                      Answer (a)

Ex. The ratio mass of oxygen and nitrogen particular gaseous mixture is 1: 4. The ratio of number of their molecule is                                                                                              (2014 Main)

 

Solution: no2nN2= (mO2)(MO2)(mN2)(MN2)  

Where, mO2  = given mass O₂, mN2  = given mass of N₂, MO2  = molecular mass of O₂, MN2  = molecular mass of N₂, nO2  = Number of moles of O₂, nN2  = number of moles of N₂

=mO2mN22832=14×2832=732                                                                                      Answer (b)

Ex.:- How many moles of electron weighs 1 kg?                                               (2002, 3M)

a 6.023×1023                                       b19.108×1031

c 6.0239.108×1054                             d19.108×6.023×108

Solution: Mass of electron = 9.108 x 10^-31 kg

9.108 x 10^-31 kg = 1.0 electron

∴   1 kg = 19.108 ×10-31  electrons= 10319.108× 16.023 ×1023  

∴  19.108×6.023 ×108 mole of electrons                                                                Answer (d)

 

MOLARITY MOLALITY, MOLE FRACTION AND NORMALITY

• Molarity is the number of moles of solute in per liter of solution. Unit is moles per liter.

Molarity= Gram Mole of solutelitre of solutiuon  

• Molality is the number of solute present in 1kg of solvent.

Molarity= Moles of solutewt. of solvent in kg  

• Normality= Gram Equivalent of soluteLitre of solution            
• Mole fraction of solute= Moles of solute Moles of solute+moles of solvent  

Ex.:  3g of activated charcoal was added to 50mL of acetic acid solution (0.06N) in a flask. After an hour it was filtered and the strength of the filtrate was found to be 0.042N. The amount of acetic acid adsorbed (per gram of charcoal) is.                                                      (2015 JEE MAINS)

(a) 18 mg                                            (b) 36 mg

(c) 42 mg                                            (d) 54 mg

Solution. Given, initial strength of acetic acid = 0.06N

Final strength = 0.042 N;        Volume = 50mL

\ Initial millimoles of CH₃COOH = 0.06 x 50 =3

Final millimoles of CH₃COOH = 0.042 x 50 = 2.1

\ Millimoles of CH₃COOH adsorbed = 3 – 2.1 = 0.9 mmol.

                                                              = 0.9 x 60mg = 54 mg                                  Answer (d)

Ex.: Dissolving 120 g of urea (mol. wt. 60) in 1000 g of water gave a solution of density 1.15 g/mL. The molarity of the solution is                                                                                        (2011)

(a) 1.78 M                                         (b) 2.00M

(c) 2.05 M                                         (d) 2.22 M

Solution: Molarity= Moles of solute Volume of Solution(L)  

Moles of urea= 12060=2  

Weight of solution = weight of solvent + weight of solute

                              = 1000 +120 = 1120 g

⇒       Volume= 1120g1.15g/mL×11000L/L=0.973L  

⇒    Molarity= 2.0000.973=2.05M                                                                                 Answer (c)

Ex.: The molarity of a solution obtained by mixing 750 mL of 0.5 M HCl 250 mL of 2 M HCl will be                                                                                                                                      (2013 Main)

(a) 0.875 M                                            (b) 1.00 M

(c) 1.75 M                                              (d) 0.0975 M

Solution: From the formula, Mf= M1V1+M2V2V1+V2  

Given, V₁ = 750 mL,  M₁ = 0.5 M

             V₂ = 250mL, M₂ = 2 M

                                                    = 750 × 0.5+250 × 2750+250= 8751000=0.875M                      Answer (a)

 

 

 

Day - 3

 

 

PERCENTAGE COMPOSITION

Mass % of an element=mass of that element in the compound ×100molar mass of the compound  

Let’s take example of carbon-dioxide (CO₂)

Mass % of C=1244 x 100=27.2%         

Mass % of O= 3244x 100=72.73%  

Ex.:- The most abundant element by mass in the body of a healthy human adult are oxygen (61.4%) carbon (22.9%), hydrogen (10.0%) and nitrogen (2.6%) The weight which a 75 kg person would gain if all Hatoms are replaced by 2 Hatoms are replaced by 2H atoms is       2017 JEE Main)

(a) 15Kg                                                  (b) 37.5 kg

(b) 7.5 Kg                                                (d) 10kg

Solution: Given, abundance of element by mass

Oxygen = 61.4% carbon= 22.9%, hydrogen = 10% and nitrogen = 2.6%

Total weight of person = 75 kg

Mass due to 1H= 75 ×10100=7.5kg    

1H atoms are replaced by 2H atoms  

Mass due to 2H=7.5 ×2kg  

\    Mass gain by person = 7.5 kg                                                                             Answer (c)

Ex.:- Given that the abundances of isotopes ₅₄Fe, ₅₆Fe and ₅₇Fe are 5% 90% and 5% respectively the atomic mass of Fe is                                                                                                            (2009)  

(a) 55. 85                                                          (b) 55.95

(c) 55.75                                                           (d) 56.05

Solution: From the given relative abundance the average weight of Fe can be calculated as

A= 54 × 5 × 56 × 90 ×57 ×5100=55.95                                                            Answer (b)

If we know percent composition we can find empirical formula

 

An empirical formula represents the simplest whole number ratio of various atoms present in a compound

•  Molecular formula shows the exact number of different types of atoms present in a molecule of a compound.
•  If the mass per cent of various elements present in a compound is known, its empirical formula can be determined.
•  Molecular formula = n x (Empirical formula), where n is a simple number and may have values 1, 2, 3….

Following steps should be followed to determine empirical formula of the compound

5.  Step 1: Conversion of mass per cent of various elements into grams.
5.  Step 2: Convert mass obtained in step1 into number of moles
5.  Step 3: Divide the mole value obtained in step 2 by the smallest mole value (out of the mole value of various elements calculated)
5.  Step 4: In case the ratios are not whole numbers, then they may be converted into whole number by multiplying by the suitable coefficient.
5.  Step 5: Write empirical formula by mentioning the numbers after writing the symbols of respective elements

Example: - In an organic compound contains 37% carbon 50% oxygen and 13% hydrogen what is the empirical formula of compound. If it’s V. D is to find the molecular formula.

Empirical formula CH₄O.

Mol. Wt. = VD x 2 = 32

Emp. formula wt = 32

n = M.W/E.W =1

So molecular formula = CH₄O

 RUTHERFORD MODEL

Observations from a - particles scattering experiment by Rutherford:

a. Most of the a - particles (nearly 99 %) passed through gold foil undeflected

b. A small fraction of a- particles got deflected through small angles

c. Very few a - particles did not pass through foil but suffered large deflection nearly 180^O

Inferences Rutherford drew from  - particles scattering experiment:

a. Since most of the a -particles passed through foil undeflected, it means most of the space in atom is empty

b. Since some of the a -particles are deflected to certain angles; it means that there is positively mass present in atom

c. Since only some of the a -particles suffered large deflections; the positively charged mass must be occupying very small space

d. Strong deflections or even bouncing back of a -particles from metal foil were due to direct collision with positively charged mass in atom

As per Rutherford the size of nucleus is 10⁴-10⁵ times smaller than atom. Mardsen & Rutherford calculated density of nucleus.

Volume of atom / Volume of nucleus = 10¹⁵

Volume of earth approximately 1.08x10²¹ cubic meters

If earth was an atom then volume of nucleus = 1.08x10²¹/10¹⁵ = 1.08x10⁶

That is ball of radius = 44m

Rutherford and Marsden gave an equation to find radius of nucleus of any atom.

R = R_O A^1/3

Ro = 1.33x 10^-13 cm

A = no. of nucleons (Neutrons + protons)

Example calculate the density of nucleus of oxygen atom.

 

 

   =1.685 x 10¹⁴ gm. /cm³ = 1.685 x 10¹⁷ kg/m³

Mass of moon = 7.34767309 × 10²² kilograms

If moon is all made of nucleus it will become a ball of radius 32m. Slightly bigger than basketball court!!!!!!!

Drawback of Rutherford model

1. According to Rutherford’s model of atom, electrons which are negatively charged particles revolve around the nucleus in fixed orbits. Thus, the electrons undergo acceleration. According to electromagnetic theory of Maxwell, a charged particle undergoing acceleration should emit electromagnetic radiation. Thus, an electron in an orbit should emit radiation. Thus, the orbit should shrink. But this does not happen.

   2. The model does not give any information about how electrons are distributed around nucleus and what are energies of these electrons

Atomic number (Z): It is equal to the number of protons in an atom. It is also equal to the number of electrons in a neutral atom.

Mass number (A): It is equal to the sum of protons and neutrons.

Isotopes: These are the atoms of the same element having the same atomic number but different mass number. Ex ₆C¹² and ₆C¹⁴

Isobars: Isobars are the atoms of different elements having the same mass number but different atomic number. ₁₈Ar⁴⁰ and ₂₀Ca⁴⁰

Isoelectronic species: These are those species which have the same number of electrons. Ex. N₂ and CO

Isotones: Two nuclides are called isotones if they have same neutron number N, but different proton numbers Z. For example, Boron- 12 and Carbon-13 nuclei, but both contain 7 neutrons

ELECTROMAGNETIC WAVES

The radiations which are associated with electrical and magnetic fields are called electromagnetic radiations. When an electrically charged particle moves under acceleration, alternating electrical and magnetic fields are produced and transmitted. These fields are transmitted in the form of waves. These waves are called electromagnetic waves or electromagnetic radiations.

PROPERTIES OF ELECTROMAGNETIC RADIATIONS:

a. Oscillating electric and magnetic field are produced by oscillating charged particles. These fields are perpendicular to each other and both are perpendicular to the direction of propagation of the wave.

b. They do not need a medium to travel. That means they can even travel in vacuum.

WAVE NATURE OF ELECTROMAGNETIC RADIATION

Characteristics of electromagnetic radiations:

a. Wavelength: It may be defined as the distance between two neighbouring crests or troughs of wave as shown. It is denoted by l.

b. Frequency (n): It may be defined as the number of waves which pass through a particular point in one second.

c. Velocity (v): It is defined as the distance travelled by a wave in one second. In vacuum all types of electromagnetic radiations travel with the same velocity. Its value is 3 X10⁸ m sec^-1. It is denoted by v

d. Wave number: Wave number (ν  ) is defined as the number of wavelengths per unit length.

Relationship between velocity, frequency and wavelength Velocity = frequency x wavelength

Velocity = frequency x wavelength

THE SPECTRUM OF ELECTROMAGNETIC RADIATION

Ex.-The Vividh Bharati station of All India Radio, Delhi, broadcasts on a frequency of 1,368 kHz (kilo hertz).

Calculate the wavelength of the electromagnetic radiation emitted by transmitter. Which part of the electromagnetic spectrum does it belong to?

        

        = 219.3m

This is a characteristic radio wave wavelength.

Ex.: Calculate (a) wave number and (b) frequency of yellow radiation having wavelength 5800 Å.

Solution

1. Calculation of wavenumber

      λ = 5800Ǻ = 5800 x 10^-8 cm

                       = 5800 x 10^-10 m

      = 1.724 x 10⁶ m^-1

      = 1.724 x 10⁴ cm^-1

(b) Calculation of the frequency (n)

TYPES OF SPECTRUM (CAN BE CLASSIFIED IN TWO DIFFERENT WAYS

FIRST CLASSIFICATION SPECTRUM IS OF TWO TYPES: CONTINUOUS AND LINE SPECTRUM

a. The spectrum which consists of all the wavelengths is called continuous spectrum.

b. A spectrum in which only specific wavelengths are present is known as a line spectrum. It has bright lines with dark spaces between them.

Electromagnetic spectrum is a continuous spectrum. It consists of a range of electromagnetic radiations arranged in the order of increasing wavelengths or decreasing frequencies. It extends from radio waves to gamma rays.

SPECTRUM IS ALSO CLASSIFIED AS EMISSION AND ABSORPTION SPECTRUM.

• Emission spectrum: A substance absorbs energy and moves to a higher energy state. The atoms, molecules or ions that have absorbed radiation are said to be excited. Since the higher energy state is unstable they return to the more stable energy state by emitting the absorbed radiation in various regions of electromagnetic spectrum. The spectrum of radiation emitted by a substance that has absorbed energy is called an emission spectrum.
• Absorption spectrum is the spectrum obtained when radiation is passed through a sample of material. The sample absorbs radiation of certain wavelengths. The wavelengths which are absorbed are missing and come as dark lines

 

BOHR’S MODEL FOR HYDROGEN ATOM

a. An electron in the hydrogen atom can move around the nucleus in a circular path of fixed radius and energy. These paths are called orbits or energy levels. These orbits are arranged concentrically around the nucleus.

b. As long as an electron remains in a particular orbit, it does not lose or gain energy and its energy remains constant.

c. When transition occurs between two stationary states that differ in energy, the frequency of the radiation absorbed or emitted can be calculated. n = DE/h = (E₂-E₁)/h

n= Frequency of radiation

h = Planck's constant

E₁ Energy of lower energy state

E₂ Energy of higher energy state

d. An electron can move only in those orbits for which its angular momentum is an integral multiple of h/2p      Þ m_evr = n   where n = 1, 2 ,3 ….

Bohr’s theory for hydrogen atom:

a. Stationary states for electron are numbered in terms of Principal Quantum numbered as n=1, 2, 3…

b. For hydrogen atom: The radii of the stationary states is expressed as r_n = n²a₀ where a₀= 52.9 pm

c. Energy of stationary state E_n = -R_H  where n 1, 2,3,.... and R_H is 2.18x10^-18 ( rhydberg’s constant)

 CALCULATION OF WAVELENGTH FOR HYDROGEN SPECTRUM

 BOHR MODEL OF THE HYDROGEN ATOM

 

 

For equilibrium Þ centripetal force = centrifugal force

 

q₁ = Ze   q₂ = e

 

On solving equation (1) & (3)

 

 

From equation (3)

Þ  r_n = r₁ x n²                   Different orbit Some Atom (DOSA)

                     Same Orbit Different Atom (SODA)

From equation (4)

Þ  V_n = V₁ / n                   DOSA

Þ V_Z   = V_H x Z                SODA

Total Energy = K.E. + P.E.

 

Putting value form equation (3) & (4)

 

From equation (5)

                             D.O.S.A

                          S.O.D.A

For Hydrogen Atom Z    =1

As per Bohr model

∆E = hv

 

 

 

R_H=109677 cm^-1 and is called Rydberg’s constant

SUCCESS AND FAILURES OF BOHR’S MODEL OF ATOM

Success of Bohr’s model

1. The frequencies of spectrum in hydrogen lines were successfully explained.

2. The value of Rydberg’s constant calculated matched the experimental value.

3. It could explain the spectrum of hydrogen like species like He⁺, Li⁺² etc

LIMITATIONS OF BOHR’S MODEL OF ATOM

1. 1. Bohr’s model failed to account for the finer details of the hydrogen spectrum. For instance splitting of a line in the spectrum into two closely spaced lines.
   2. Bohr’s model was also unable to explain spectrum of atoms containing more than one electron.
   3. Bohr’s model was unable to explain Zeeman Effect i.e. splitting of spectral line in presence of magnetic effect.
   4. Bohr’s model also failed to explain Stark effect i.e. splitting of spectral line in presence of electric field.
   5. Bohr’s model could not explain the ability of atoms to form molecules by chemical bonds

Black body: An ideal body, which emits and absorbs all frequencies, is called a black body. The radiation emitted by such a body is called black body radiation.

Planck’s Quantum Theory

Planck said that atoms and molecules could emit (or absorb) energy only in discrete quantities and not in a continuous manner, a belief popular at that time.

Planck gave the name quantum to the smallest quantity of energy that can be emitted or absorbed in the form of electromagnetic radiation.

E µ n    Þ  E = hn Þ

The proportionality constant, ’h’ is known as Planck’s constant and has the value 6.626 10^-34 Js.     

PHOTOELECTRIC EFFECT

Experimental results observed for the experiment of Photoelectric effect observed by Hertz:

a. When beam of light falls on a metal surface electrons are ejected immediately i.e. there is not time lag between light striking metal surface and ejection of electrons

b. Number of electrons ejected is proportional to intensity or brightness of light

c. Threshold frequency ( n_o ): For each metal there is a characteristic minimum frequency below which photoelectric effect is not observed. This is called threshold frequency.

d. If frequency of light is less than the threshold frequency there is no ejection of electrons no matter how long it falls on surface or how high is its intensity.

Photoelectric work function (W_o): The minimum energy required to eject electrons is called photoelectric work function. W_o= h n_o

Energy of the ejected electrons: ½ m_ev² = h(n-n_o)

 The kinetic energies of these electrons increase with the increase of frequency of the light used. All the above results could not be explained on the basis of laws of classical physics.

Photoelectric work function (W_o): The minimum energy required to eject electrons is called photoelectric work function. W_o= h n_o

Energy of the ejected electrons: ½ m_ev² = h(n-n_o)

DUAL BEHAVIOUR OF ELECTROMAGNETIC RADIATION

It is concluded from active experiment that electromagnetic radiation have due nature. Wave nature as confirmed by experiments and particle nature is confirmed by photo electric effect.

de Broglie proposed that matter exhibits dual behaviour i.e. matter shows both particle and wave nature.

1. de Broglie’s relation: l = =

Where: l - Wavelength;  p – Momentum;   v – Velocity;  h – Planck’s constant

2. According to de Broglie, every object in motion has a wave character. Wavelengths of macroscopic objects cannot be detected but for microscopic particles it can be detected. This is because for microscopic objects, the mass is less. Since mass and wavelength are inversely proportional to each other, the wavelength will be more. But for macroscopic objects, the mass is large. Therefore, wavelength will be too short to be detected. Dual

Heisenberg’s uncertainty principle

It states that it is impossible to determine simultaneously, the exact position and exact momentum (or velocity) of an electron

 

Assuming mass is constant and does not change with velocity

Schrodinger wave equation

The probability of finding an electron at a point within an atom is proportional to the  square of the orbital wave function i.e., ψ² at that point. ψ² is known as probability density and is always positive. From the value of ψ2 at different points within an atom, it is possible to predict the region around the nucleus where electron will most probably be found

IMPORTANT FEATURES OF THE QUANTUM MECHANICAL MODEL OF ATOM

Quantum mechanical model of atom is the picture of the structure of the atom, which emerges from the application of the Schrödinger equation to atoms. The following are the important features of the quantum me chemical model of atom:

1. The energy of electrons in atoms is quantized (i.e., can only have certain specific values), for example when electrons are bound to the nucleus in atoms.

2. The existence of quantized electronic energy levels is a direct result of the wave like properties of electrons and is allowed solutions of Schrödinger wave equation.

3. Both the exact position and exact velocity of an electron in an atom cannot be determined simultaneously (Heisenberg uncertainty principle). The path of an electron in an atom therefore, can never be determined or known accurately. That is why, as you shall see later on, one talks of only probability of finding the electron at different points in an atom.

4. An atomic orbital is the wave function ψ for an electron in an atom. Whenever an electron is described by a wave function, we say that the electron occupies that orbital. Since many such wave functions are possible for an electron, there are many atomic orbitals in an atom. These one electron orbital wave functions.  or orbitals form the basis of the electronic structure of atoms. In each orbital, the electron has a definite energy. An orbital cannot contain more than two electrons. In a multi-electron atom, the electrons are filled in various orbitals in the order of increasing energy. For each electron of a multi-electron atom, there shall, therefore, be an orbital wave function characteristic of the orbital it occupies. All the information about the electron in an atom is stored in its orbital wave function ψ and quantum mechanics makes it possible to extract this information out of ψ.

ORBITALS AND QUANTUM NUMBERS

ORBIT VS. ORBITAL

 PHYSICAL SIGNIFICANCE OF THE QUANTUM NUMBERS

1s and 2s orbitals

 2s and 2p Orbitals

d orbital

 ELECTRONIC CONFIGURATION OF ATOMS

Aufbau principle Electrons are placed in the lowest energetically available subshell. Thus orbitals are filled in the order of increasing energy, using two general rules  to help predict electronic configurations:-

1. Electrons are assigned to orbitals in order of increasing value of (n+ℓ). Orbitals with lower value of (n+l) are filled first as they have lower energy.
2. For subshells with the same value of (n+ℓ), electrons are assigned first to the sub shell with lower n.
3. The order in which the orbitals are filled is as follows:
4. 1s, 2s, 2p, 3s, 3p, 4s, 3d, 4p, 5s, 4d, 5p, 4f, 5d, 6p, 7s...
5. It is based on (n+ l) rule. It states that the orbital with lower value of (n +l) has lower energy.

 AUFBAU DIAGRAM

PAULI'S EXCLUSION PRINCIPLE

It is impossible for any two electrons of an atom to have the same values of the all four quantum numbers. If even three quantum numbers are same for any two electrons of the atom then fourth quantum number will definitely be different.

In other words an orbital can have maximum of two electrons

Hund’s Rule of Maximum Multiplicity

This rule deals filling of electrons in the orbitals belonging to the same subshells of equal energy called degenerate orbitals.

5.  It states that pairing of electrons in the orbitals belonging to the same sub shell (p, d or f) does not take place until each orbital of that sub shell gets one electron that is singly occupied.
5.  Some of the orbitals acquire extra stability due to their symmetry.

 EXCEPTION: HALF-FILLED AND FILLED D ORBITALS

•  When a subshell is half-filled or completely filled, it lowers the overall energy of the atom.
•  In other words, this is favoured.
•  This usually doesn't affect electron configurations, except in the d and f subshells.
•  Example: Copper (29 electrons)
•  Normal electron configuration:

      1s²2s²2p⁶3s²3p⁶4s²3d⁹

•  Added stability if 3 d is full (electron gets promoted to 3 d from 4s)
•  1s²2s²2p⁶3s²3p⁶4s¹3d¹⁰

The region where this probability density function reduces to zero is called nodal surfaces or simply nodes.

Boundary surface diagram: In this representation, a boundary surface or contour surface is drawn in space for an orbital on which the value of probability density ψ ²(r) is constant. However, for a given orbital, only that boundary surface diagram of constant probability density is taken to be good representation of the shape of the orbital which encloses a region or volume in which the probability of finding the electron is very high, say, 90%.

Radial nodes: Radial nodes occur when the probability density wave function for the electron is zero on a spherical surface of a particular radius. Number of radial nodes = n – l – 1

Angular nodes: Angular nodes occur when the probability density wave function for the electron is zero along the directions specified by a particular angle. Number of angular nodes = l

Total number of nodes = n – 1

Degenerate orbitals: Orbitals having the same energy are called degenerate orbitals

IMPORTANT FORMULAE

1.         
2. Frequency of electromagnetic wave  
3. Energy of photon E = hn =  
4. 
5. Bohr’s postulate mvr =  (Angular momentum of ‘e’ in an orbit)            
6. 
7. r_n = r₁ x n²  same atom different orbit
8. r_Z  = r_H / Z   same orbit different atom
9. 
10. u_n = u₁ / n        same atom different orbit
11. u_Z = u_H x Z      same orbit different atom
12.   E₁ for H= -21.72X10^-12 erg  = -13.6 eV
13. E_n = E₁/n²             same atom different orbit
14. E_Z = E_H x Z²    same orbit different atom
15. 
16. A^o
17. Ön=a (Z – b), where n is frequency of X-rays given out by metal of at. No. Z
18. Average atomic weight =  
19. KE of photoelectric electron (½ mv²) = hn - hn_o (hn_o = work function) & n_o threshold frequency
20. deBroglie equation l =  
21. Uncertainty
22. (Schrodinger wave equation)  
23. Angular momentum of ‘e’ in an orbital=  
24. Magnetic moment of an atom=.  where n is no. of unpaired electrons
25. Nodes
    1. Radial (spherical) nodes=(n – l – 1)
    2. Angular (non spherical) nodes=l
    3. Total nodes= (n - 1) 
26. No of spectral lines =  

LAWS OF CHEMICAL COMBINATIONS

There are 5 basic laws of chemical combinations that govern every reaction: Law of conservation of mass, law of definite proportions, law of multiple proportions, Gay Lussac’s law of gaseous volumes, and lastly, Avogadro law

• Law of Conservation of Mass: Antoine Lavoisier established the Law of Conservation of Mass. It states that matter can neither be created nor destroyed. In other words, we can say that during any physical or chemical change, the total mass of reactants is equal to the total mass of products
• Law of definite proportions: Joseph Proust showed that a given compound always contains exactly the same proportion of elements by weight.
• Law of multiple proportions: Dalton proposed the law of multiple proportions. According to this law if two elements can combine to form more than one compound, the mass of one element that combines with the fixed mass of the other element is in the ratio of small whole numbers.
• Gay Lussac’s Law of gaseous volumes: When gases combine or are produced in a chemical reaction they do so in a simple ratio by volume, provided all the gases are at same temperature and Pressure.
• Avogadro law: At the same temperature and pressure, equal volumes of gases contain equal number of molecules.  

INTRODUCTION

In 1800, only 31 elements were known. By 1865, the number of identified elements had more than doubled to 63. At present 114 elements are known.

With such a large number of elements it is very difficult to study individually. the chemistry of all these elements is also difficult to handle individually.

LAW OF TRIADS

Johann Döbereiner classified elements in group of three elements called triads.

In Döbereiner triad the atomic weight of the middle element is very close to the arithmetic mean of the other two elements

LAW OF OCTAVES

Since Döbereiner Law of triads worked only for few elements, it was dismissed.

Chancourtois arranged elements in order of increasing atomic weights and made a cylindrical table of elements.

John Newland arranged the elements in the increasing order of atomic weight and noted that the properties of the every eighth element are similar to the first one. This relationship is called as “Law of octaves”

MENDELEEV’S PERIODIC TABLE

According to Mendeleev’s periodic law the physical and chemical properties of elements are periodic functions of their atomic weights.

Merits of Mendeleev’s periodic table:

• Mendeleev’s periodic table was very helpful in remembering and studying the properties of large number of elements
• Mendeleev’s periodic table helped in correcting the atomic masses of some of the elements like gold, beryllium and platinum based on their positions in the periodic table
• Mendeleev could predict the properties of some undiscovered elements like scandium, gallium and germanium. By this intuition, he had left gaps for the undiscovered elements while arranging elements in his periodic table

 SUCCESS OF MENDELEEV’S WORK

Limitation of Mendeleev’s table

• Position of hydrogen is not correctly defined in periodic table. It is placed in group I though it resembles both group 1 and 17.
• In certain pairs of elements increasing order of atomic masses was not obeyed. For example argon (Ar, atomic mass 39.9) is placed before potassium (K, atomic mass 39.1)
• Isotopes were not given separate places in the periodic table although Mendeleev's classification is based on the atomic masses.
• Some similar elements are separated and dissimilar elements are grouped together. For example copper and mercury resembled in their properties but had been placed in different groups. On the other hand lithium and copper were placed together although their properties are quite different.
• Mendeleev did not explain the cause of periodicity among the elements.

• Lanthanoids and actinoids were not given a separate position in the table.

Modern Periodic Table

Henry Moseley showed that the atomic number is a more fundamental property of an element than its atomic mass.

Mendeleev’s Periodic Law was, therefore, accordingly modified.

Modern Periodic Law:

The physical and chemical properties of the elements are periodic functions of their atomic numbers

 Nomenclature of elements of Atomic mass > 100

Electronic Configurations for Periods and groups

Electronic configuration of Zirconium (A. No. 40): 1s² 2s² 2p⁶ 3s² 3p⁶ 4s² 3d¹⁰ 4p⁶ 5s² 4d²

Iodine I₅₃  = 1s² 2s² 2p⁶ 3s² 3p⁶ 4s² 3d¹⁰ 4p⁶ 5s² 4d¹⁰ 5p⁵

We can classify the elements into four blocks viz.,

s-block, p-block, d-block and f-block depending on the type of atomic orbitals that are being filled with electrons.

Hydrogen and Helium are exception to this rule

The s-Block Elements

The elements of Group 1 (alkali metals) and Group 2 (alkaline earth metals) which have ns¹ and ns² outermost electronic configuration belong to the s-Block Elements.

The p-Block Elements

The outermost electronic configuration varies from ns²np¹ to ns²np⁶ in each period. The p-Block Elements comprise those belonging to Group 13 to 18 and these together with the s-Block Elements are called the Representative Elements or Main Group Elements.

The d-Block Elements (Transition Elements)

These elements have the general outer electronic configuration (n-1)d^1-10ns^0-2 .  These are the elements of Group 3 to 12 in the centre of the Periodic Table. These are characterised by the filling of inner d orbitals by electrons and are therefore referred to as d-Block Elements. They are all metals.

The f-Block Elements (Inner-Transition Elements)

They have outer electronic configuration (n-2)f^1-14 (n-1)d^0-1ns²

The two rows of elements at the bottom of the Periodic Table, called the  Lanthanoids, Ce(Z = 58) to Lu(Z = 71) and Actinoids, Th(Z = 90) to Lr (Z = 103) .

METALS, NON METALS AND METALLOIDS

Metals comprise more than 78% of all known elements and appear on the left side of the Periodic Table. Metals are usually solids at room temperature [mercury is an exception; gallium and caesium also have very low melting points (303K and 302K, respectively)]. Metals usually have high melting and boiling points. They are good conductors of heat and electricity. They are malleable (can be flattened into thin sheets by hammering) and ductile (can be drawn into wires). In contrast, non-metals are located at the top right hand side of the Periodic Table. In fact, in a horizontal row, the property of elements change from metallic on the left to non-metallic on the right. Non-metals are   usually solids or gases at room temperature with low melting and boiling points (boron and carbon are exceptions). They are poor conductors of heat and electricity.

Most non-metallic solids are brittle and are neither malleable nor ductile. The elements become more metallic as we go down a group; the non-metallic character increases as one goes from left to right across the Periodic Table. The change from metallic to non-metallic character is not abrupt as shown by the thick zigzag line in periodic table.

The elements (e.g., silicon, germanium, arsenic, antimony and tellurium) bordering this line and running diagonally across the Periodic Table show properties that are characteristic of both metals and non-metals. These elements are called Semi-metals or Metalloids

SHIELDING EFFECT

Shielding effect or screening effect: Due to the presence of electrons in the inner shells, the electron in the outer shell will not experience the full positive charge on the nucleus.

So due to the screening effect, the net positive charge experienced by the electron from the nucleus is lowered and is known as effective nuclear charge.

• Effective nuclear charge, Z_eff, experienced by an electron is less than actual nuclear charge , Z
• Electrons in the outermost shell are repelled (shielded) by electrons in the inner shells. This repulsion counteracts the attraction caused by the positive nuclear charge

Coulomb’s law:

 SHIELDING EFFECT

Effective nuclear charge, Z_eff, experienced by an electron is less than the actual nuclear charge, Z. „Electrons in the outermost shell are repelled (shielded) by electrons in the inner shells. This repulsion counteracts the attraction caused by the positive nuclear charge

Z_eff = Z – S (S = screening constant)

Shielding effect

1. Electrons in inner orbitals have greater shielding effect than electrons in same shell.
2. Shielding effect s > p > d > f

ATOMIC RADII

Periodicity

• As we move from left to right along the period, the effective nuclear charge "felt" by the outermost electron increases while the distance from the nucleus doesn't change that much (electrons are filling the same shell)
• Outermost electrons are attracted stronger by the nucleus, and the atomic radius decreases

IONIC RADII

When atom loses an electron, its radius always decreases

• Cations (positive ions) are always smaller than their respective neutral atoms.

When atom gains an electron, its radius always increases

• Anions (negative ions) are always larger than their  respective neutral atoms

 ISOELECTRONIC SPECIES

RADII OF ISOELECTRONIC IONS

IONIZATION ENERGY

• If sufficient energy is provided, the attraction between the outer electron and the nucleus can be overcome and the electron will be removed from the atom

FIRST IONIZATION ENERGY (IE₁)

• The minimum amount of energy required to remove the most loosely bound electron from an isolated gaseous atom to form a 1⁺ ion

Na (g) + 496 KJ/ Mol → Na⁺ (g) + e^-

• Second ionization energy (IE₂)

The minimum amount of energy required to remove the 2nd electron from a gaseous 1+ ion

IE₁:    Ca(g) + 590 kJ/mol ® Ca⁺(g) + e^-

IE₂:  Ca⁺(g) + 1145 kJ/mol  ® Ca²⁺(g) + e^-

The 2nd electron "feels" higher nuclear charge (stronger attractive force) since the electron-electron repulsion has been decreased: For all atoms IE₂ > IE₁ 

Ionization Energy: Trends

• IE₁ increases from left to right along a period since the effective nuclear charge (Z_eff)  "felt" by the outermost electrons increases

• There are some exceptions to this general trend caused by additional stability of filled and half- filled subshells (orbitals with the same )
• IE₁ decreases as we go down a group since the outermost electrons are farther from the nucleus

IONIZATION ENERGY: PERIODICITY

IMPORTANT CONCLUSIONS

• Atoms of noble gases have completely filled outer shell, the smallest radii among the elements in the same period, and the highest ionization energies
• Atoms of metals, especially those to the left in the periodic chart, ionize easily forming cations and attaining the electron configuration of noble gases
• Atoms of non-metals, especially those to the right in the periodic chart, are very unlikely to loose electrons easily - their ionization energies are high

ELECTRON AFFINITY

• For most non-metals, it is much easier to achieve the stable electron configuration of a noble gas by gaining rather than losing electrons
• Therefore, non-metals tend to form anions
•  Electron affinity is a measure of an atom's ability to form negative ions
• The amount of energy released when an electron is added to an isolated gaseous atom to form an ion with a -1 charge.

                                                      Cl(g) + e^-   ® Cl^-(g) + 349 kJ/mol

Sign conventions for electron affinity

• If electron affinity < 0 (energy is released)

• If electron affinity > 0 (energy is absorbed)

Compare cation- and anion-forming processes:

ELECTRON AFFINITY: TRENDS

EA becomes more negative on going from left to right along a period

• There are some exceptions to this general trend caused by

additional stability of filled and half- filled subshells

(orbitals with the same l )

• EA becomes less negative as we go down a group because

the attraction of the outermost electrons to the nucleus weakens

ELECTRON AFFINITY: PERIODICITY

• Important conclusions

• Noble gases have completely filled outer shell and therefore zero electron affinity
• Non-metals, especially halogens, gain electrons easily forming anions and attaining the electron configuration of noble gases
• Metals are usually quite unlikely to gain electrons and form anions

Electron Gain Enthalpies (kJ mol^-1) of group17 elements

 

ELECTRONEGATIVITY

• Measures the tendency of an atom to attract electrons when chemically combined with another element

• If element "likes" electrons - high electronegativity (electronegative element)

• If element "dislikes" electrons - low electronegativity (electropositive element)

Sounds like the electron affinity but different

• Electron affinity measures the degree of attraction of an electron by a single atom forming an anion

• Electronegativity measures the attraction of electrons to the atom in chemical compounds

• The scale for electronegativity was suggested by Linus Pauling
• It is a semi-qualitative scale based on data collected from studying many compounds ​​​​​​​ 

ANOMALOUS PROPERTIES OF SECOND PERIOD ELEMENTS

The first element of each of the groups 1 (lithium) and 2 (beryllium) and groups 13-17 (boron to fluorine) differs in many respects from the other members of their respective group. Behaviour of lithium and beryllium is more similar with the second element of the following group i.e., magnesium and aluminium, respectively. This sort of similarity is commonly referred to as diagonal relationship in the periodic properties. What are the reasons for the different chemical behaviour of the first member of a group of elements in the s- and p-blocks compared to that of the subsequent members in the same group? The anomalous behaviour is attributed to their small size, large charge/ radius ratio, high Electronegativity and absence of d orbital

INTRODUCTION

The attractive force which holds together the constituent particles (atoms, ions or molecules) in chemical species is known as chemical bond.

1. Kössel-Lewis Approach to Chemical Bonding

They assumed that atom have positive kernel surrounded by electrons occupying the corners of a cube. If they have all the eight electrons in their outer shell they will be stable (octet rule). Otherwise they achieve stability (octet) through chemical bonding.

     LEWIS STRUCTURE

Why chemical bonds are formed.

If the resultant molecule has lower Gibb's energy then the reacting species, then chemical bonds are formed.

IONIC BOND

Ionic bonds are strong electrostatic forces between cation & anion which are formed when an atom looses an electron or gains an electron.

 

BORN HABER CYCLE FOR SODIUM CHLORIDE

    

Ex.1:  Draw Born Haber cycle for formation of Magnesium chloride

 

PROPERTIES OF IONIC COMPOUNDS

1. They are crystalline in nature
2. They have high Melting and Boiling point
3. Hard and Brittle
4. Soluble in polar solvents
5. Conduct electricity in molten state and aqueous state but not in solid state.
6. Do not show Isomerism.

VARIABLE ELECTROVALENCY

Fe (26) → 3s² 3p⁶ 3d⁶ 4s²

Fe⁺² (24) → 3s² 3p⁶ 3d⁶ (less stable)

Fe⁺³ (23) → 3s² 3p⁶ 3d⁵ (more stable)

Solubility of ionic compounds in water. There are two things happen on dissolving ionic compounds in water.

1. Breaking of ionic lattice

2. Mixing of ions in water

For the first process, lattice energy to be provided to the solution and for second process hydration energy will be released by the system

If ΔH_hyd >ΔH_L.E. Compound is soluble in water

(a) Lattice Energy

1. If depends on size of cation and anion.

Smaller the size greater is the Lattice Energy.

Ex. → LiCl > NaCl > KCl

      → NaCl > NaBr > NaI

2. If depends on Charge of cation and anion.

Bigger the charge, lesser is it Lattice energy.

Ex. MgCl₂ > NaCl

3. In case anions are extremely large, then the rule one charges smaller the cation lesser is Lattice Energy.

Ex. MgSO₄ < CaSO₄ < SrSO₃ < BaSO₄

(b) Hydrogen energy

Smaller the cation, greater is the hydration enthalpy.

Ex. Compare Solubility BaSO₄ and CaSO₄

L.E. Þ BaSO₄ > CaSO₄

Hyd. Energy Þ CaSO₄ > BaSO₄

So CaSO₄ is more soluble.

Ex. Compare solubility of NaCl and BaCl₂

L.E. Þ BaCl₂ > NaCl

Hyd. Energy Þ NaCl > BaCl₂

So NaCl is more soluble.

 

COVALENT BOND

Combining of unpaired electrons of atoms to achieve a stable configuration and formation of molecules is called covalent bond.

SIGMA (s) BOND: THIS TYPE OF COVALENT BOND is formed by the end to end (head-on) overlap of bonding orbitals along the inter-nuclear axis. This is called as head on overlap or axial overlap

s - s overlapping: In this case, there is overlap of two half killed s-orbitals along the inter-nuclear axis as shown below:

 

s-p overlapping: This type of overlap occurs between half filled s-orbitals of one atom and half filled p-orbitals of another atom:

 

p-p overlapping: This type of overlap takes place between half filled p-orbitals of the two approaching atoms:

 

Pi (π) BOND: IN THE FORMATION OF π BOND

The atomic orbitals overlap in such a way that their axes remain parallel to each other and perpendicular to the inter-nuclear axis.

 

Characteristic of Convent compounds

1. Physical state ® gases, liquids of law b.p. & soft solids
2. Melting Point/Boiling point ®  with exception of network solids, they have low M.P. and B.P.
3. Electrical conductance ®  generally bad conductors
4. Solubility  ®  soluble in non-polar solvents
5. Isomerism  ®  yes Physical state ® gases, liquids of law b.p. & soft solids
6. Melting Point/Boiling point ®  with exception of network solids, they have low MP/B.
7. Electrical conduce ®  generally bed conductors
8. Solubility  ®  soluble in non-polar solvents
9. Isomerises  ®  yes

COMPARISON BETWEEN IONIC AND COVALENT BONDS

COMPARISON BETWEEN IONIC AND COVALENT COMPOUNDS

Some representative Bond Energies and Bond Lengths

COORDINATE BOND

Dative bond or semi-polar bond.

METALLIC BOND

The bonding which holds metals atoms firmly together on account of force & attraction between metal cation and mobile sea electrons is called metallic bonding.

PROPERTIES OF METALS & NON METALS

LANDON FORCES AND DISPERSION FORCES

London forces between two Helium atoms

DIPOLE - DIPOLE FORCES

Dipole-dipole forces act between the molecules possessing permanent dipole. Ends of the dipoles  possess partial charges, and these charges are shown by Greek letter delta (δ). Partial charges are always less than the unit electronic charge. The polar molecules interact with neighboring molecules.

Dipole-dipole interaction energy between stationary polar molecules (as in solids) is proportional to ,  Where r is the distance between polar molecules.

DIPOLE INDUCED DIPOLE FORCES

This type of attractive forces operate between the polar molecules having permanent dipole and the molecules lacking permanent dipole.

Permanent dipole of the polar molecule induces dipole on the electrically neutral molecule by deforming its electronic cloud

 

HYDROGEN BOND

This is special case of dipole-dipole interaction.. This is found in the molecules in which highly polar N-H, O-H or H-F bonds are present. Although hydrogen bonding is regarded as being limited to N, O and F; but species such as Chlorine may also participate in hydrogen bonding. Energy of hydrogen bond varies between 10 to 100 kJ mol.

Hydrogen bond can be classified in three types

1. Inter molecular H bond. Same molecule

1. Inter molecular H bond. Different molecule

Hydrogen Bonding between Ammonia and Water

2. Intra Molecular Hyrogen Bond

 HYDROGEN BOND STRENGTH

Hydrogen bonds can vary in strength from weak (1–2 kJ mol⁻¹) to strong (161.5 kJ mol⁻¹ in the ion Typical enthalpies in vapour include:

F − H···: F (161.5 kJ/mol or 38.6 kcal/mol), illustrated uniquely by HF₂⁻,

O − H···: N (29 kJ/mol or 6.9 kcal/mol), illustrated water-ammonia

O − H···: O (21 kJ/mol or 5.0 kcal/mol), illustrated water-water, and alcohol-alcohol

N − H···: N (13 kJ/mol or 3.1 kcal/mol), illustrated by ammonia-ammonia

N − H···: O (8 kJ/mol or 1.9 kcal/mol), illustrated water-amide

Bond Length

It is defined as the angle between the orbitals containing bonding electron pairs around the central atom in a molecule/complex ion. Bond angle is always determined experimentally

A. Ionic Compound:- bond length is sum of cationic radius and anionic radius.

B. Covalent compound:

The bond length is measured by spectroscopic, x-ray diffraction and electron diffraction. Covalent radius is then calculated from this.

Lewis structure:- unable to give bond angle

VESPR they:- much better in predicate bond angle

M. O. T:- Give bond angle much better

Bond Energy or Enthalpy

It is defined as the amount of energy required to break one mole of bonds of a particular type between two atoms in a gaseous state.

For a diatomic molecular Bond energy simply energy required to break the bond.

For Hetero-nuclear molecular it is average energy or mean energy.

Ex.   H₂O(g) → H(g) + OH(g)                  ∆H = 502

OH(g) → H(g) + O(g)                              ∆H = 427

Bond length or bond distance is the average distance between nuclei of two bonded atoms in a molecule. It is a transferable property of a bond between atoms of fixed types, relatively independent of the rest of the molecule

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1. Sigma bond (s )> Pie (p) bond

2. Triple bond > double bond > single bond

3. s – s overlap > s – p overlap > p – p overlap

OCTET RULE

Kössel and Lewis in 1916 developed an important theory of chemical combination between atoms known as electronic theory of chemical bonding. According to this, atoms can combine either by transfer of valence electrons from one atom to another (gaining or losing) or by sharing of valence electrons in order to have an eight electrons in their valence shells.

Exceptions to the Octet Rule:

1. Hydrogen molecule: Hydrogen has one electron in its first energy shell (n = 1). It needs only one more electron to fill this shell, because the first shell cannot have more than two electrons. This configuration (1s²) is similar to that of noble gas helium and is stable. In this case, therefore, octet is not needed to achieve a stable configuration

Incomplete octet of the central atom: The octet rule cannot explain the formation of certain molecules of lithium, beryllium, boron, aluminium, etc.  (LiCl, BeH₂, BeCl₂, BH₃, BF₃) in which the central atom has less than eight electrons in the valence shell as shown below:

Expanded octet of the central atom: There are many stable molecules which have more than eight electrons in their valence shells. For example, PF₅, has ten; SF₆ has twelve and IF₇ ha fourteen electrons around the central atoms, P, S, and I respectively  

Odd electron molecules: There are certain molecules which have odd number of electrons, like nitric oxide, NO and Nitrogen dioxide, NO₂. In these cases, octet rule is not satisfied for all the atoms.

 It may be noted that the octet rule is based upon the chemical inertness of noble gases. However, it has been found that some noble gases (especially xenon and krypton) also combine with oxygen and fluorine to form a large number of compounds such a XeF₂, KrF₂, XeOF₂, XeOF₄, XeF₆, etc.

This theory does not account for the shape of the molecules. It cannot explain the relative stability of the molecule in terms of the energy.

LEWIS STRUCTURE

    Lewis structure and bonding theory cannot explain the shape of molecule.

Shape of H₂O.

H → 1s²

O → 1s² 2s² px² py¹ pz¹

Angle H ─ O ─ H

Is 90^o as per this mode but shell L = 104.5^o