organic and inorganic chemistry

Inorganic and Organic Chemistry
• Overview of the Periodic Table – how do you read it?
©Dr. M.S. Shaikh-2018 69

• Overview of the Periodic Table – what is atomic number?
• The number of protons found in the
nucleus of an atom.
Or
• The number of electrons surrounding the
nucleus of an atom.

• Overview of the Periodic Table – what is symbol?
An abbreviation of element name

• Overview of the Periodic Table – what is atomic weight?
The number of protons and neutrons in the nucleus of an atom

• Overview of the Periodic Table – how to find number of protons,
electrons and neutrons in an element using periodic table?
• Number of protons = Atomic number
• Number of electrons = Atomic number
• Number of neutrons = Atomic weight – atomic number

• Overview of the Periodic Table – what is an element?
• A substance composed of a single kind of atom.
• Cannot be broken down into another substance
by chemical or physical means.

• Overview of the Periodic Table – what is a compound & mixture?
• A substance in which two or more different
elements are chemically bonded together.
• Two or more substances that are mixed
together but are not chemically bonded.

• Overview of the Periodic Table – Example of element, compound
and mixture
• Sodium is an element.
• Chlorine is an element.
• When sodium and chlorine bond they make the compound sodium
chloride, commonly known as table salt.
• When water is added, it becomes mixture.
• Compounds have different properties than the
elements that make them up.
• Table salt has different properties than sodium,
an explosive metal, and chlorine, a poisonous
gas.

• Periodic Table – more details
• The periodic table organizes the elements in a particular way.
• A great deal of information about an element can be gathered from its
position in the periodic table.
• For example, you can predict with reasonably good accuracy the
physical and chemical properties of the element.
• You can also predict what other elements a particular element will react
chemically.
• Understanding the organization and plan of the periodic table will help
you obtain basic information about each of the 118 known elements.

• Periodic Table – more details – continue…
• Elements are organized on the table according to their atomic number,
usually found near the top of the square.
• The atomic number refers to how many protons an atom of that element
has.
• For instance, hydrogen has 1 proton, so it’s atomic number is 1.
• The atomic number is unique to that element.
• No two elements have the same atomic number.

• Periodic Table – what is a square in the table?
• Different periodic tables can include various bits of information, but
usually:
• Atomic number
• Symbol
• Atomic mass
• Number of valence electrons
• State of matter at room temperature.

• Periodic Table – valence electrons
• The number of valence electrons an atom has may also appear in a
square.
• Valence electrons are the electrons in the outer energy level of an
atom.
• These are the electrons that are transferred or shared when atoms
bond together.

• Periodic Table – Categories
• The elements of periodic table are divided into three categories.
• Metals
• Non Metals
• Metalloids

• Periodic Table – Categories - Metals
• Metals are good conductors of heat and electricity.
• Metals are shiny.
• Metals are ductile (can be stretched into thin wires).
• Metals are bendy (can be pounded into thin sheets).
• A chemical property of metal is its reaction
with water which results in corrosion.

• Periodic Table – Categories - Non Metals
• Non-metals are poor conductors of heat and electricity.
• Non-metals are not ductile or malleable.
• Solid non-metals are brittle and break easily.
• They are dull/ weak.
• Examples . . .

• Periodic Table – Categories - Metalloids
• Metalloids (metal-like) have properties of both metals and non-metals.
• They are solids that can be shiny or dull.
• They conduct heat and electricity better than non-metals but not as well
as metals.
• They are ductile and malleable.
• Examples…

• Periodic Table – Arrangements
• Elements in periodic table are grouped into:
• Families / groups – vertical columns
• Elements in families have similar properties.
• Periods / series – horizontal rows
• Elements in periods do not have similar properties.

• Periodic Table – Arrangements – continue…
• Columns of elements are called
groups or families.
• Elements in each family have
similar but not identical
properties.
• For example, lithium (Li), sodium
(Na), potassium (K), and other
members of family IA are all soft,
white, shiny metals.
• All elements in a family have the
same number of valence
electrons.
• Each horizontal row of elements
is called a period.
• The elements in a period are not
alike in properties.
• In fact, the properties change
greatly across given row.
• The first element in a period is
always an extremely active solid.
• The last element in a period, is
always an inactive gas.
Families Periods

Periodic Table
A lot of more information to be covered by:
• Self reading...
• Students presentations…

• Chemistry of solutions
• Few concepts…
Solution?
• Mixture of two or more substances in a single phase.
• Once constituent is regarded as solvent and the other as solute.
• Solutions are homogeneous mixtures of two or more pure substances.
• In a solution, the solute is dispersed uniformly throughout the solvent.

• Few concepts…
Solute?
• The part of a solution that is being dissolved.
• Usually the lesser in amount.

• Few concepts…
Solvent?
• The part of a solution that dissolves the solute.
• Usually the greater in amount.

• Few concepts
Solute + Solvent = Solution?
Solute
Solvent

• Few concepts…
Miscible substances?
• Liquids that will dissolve in each other.

• Few concepts…
Electrolytes?
• Liquids that are capable of conducting electricity.

• Types of solutions
• Liquid solutions - salt and water, sugar and water
• Gaseous solutions - Air and other mixture of gases
• Solid solutions - combination of solids, homogenous
mixture of metals (alloys)

• Types of solutions – concentration terms
• Concentrated solutions – solutions with lots of solute per solvent
• Dilute solutions – solutions with lots of solvent.
• Saturated solutions – solutions with maximum amount of solute
• Super saturated solutions – solution with more solute than it can
hold.

Dilute solution Concentrated solution

Super saturated solution
• A solution that contains more solute than it can theoretically
hold at a given temperature.
• If more solute is added in the solution,
the formation of crystals increases.
• This solution type is unstable.

To be continued…

• Chemistry of solutions – how solutions are formed?
• Solvent molecules attracted to surface ions.
• Each ion is surrounded by solvent molecules.
• Enthalpy changes with each interaction (broken / formed).

• The ions are solvated (surrounded by solvent).
• If the solvent is water, the ions are hydrated.
• The intermolecular forces.

• Solvation
• It is an interaction of a solute with the solvent, which leads to
stabilization of the solute species in the solution.
• In the solvated state, an ion in a solution is surrounded or
complexed by solvent molecules.

• Solvation – an example of salt and water

• Solvation – rate of solvation – the influencing parameters
• Stirring / shaking
• Breaking apart a solute
• Heating the solvent / temperature

How does a solid dissolve in a liquid?
This gives a concept of dissolution.
What is dissolution?

• Chemistry of solutions – Dissolution?
• A process by which a solid, liquid or gas forms a solution in a
solvent.
• In solids, this can be explained as the breakdown of crystal lattice
into individual ions, atoms or molecules and their transport into the
solvent.

• Chemistry of solutions – Dissolution?
• For liquids and gases, the molecules must be compatible with those
of solvent for a solution to form.
• In this process, solid substance solubilizes in a given solvent.

• Chemistry of solutions – Dissolution – example …

Dissolution versus reaction
• Dissolution is a physical change.
• Original solute can get back by evaporating the solvent in
dissolution process.
• However, it is not possible to get back original solute in a
reaction.

Solubility
• Solubility is a chemical property referring to the ability for a
given substance, the solute, to dissolve in a solvent.
• It is measured in terms of the maximum amount of solute
(expressed in grams) dissolved in a solvent (100 grams) at
equilibrium (specific temperature and pressure).
• The resulting solution is called a saturated solution.

• Chemistry of solutions – solubility
Factors affecting solubility
• Temperature
• Pressure
• Molecular size
• Stirring
• Intermolecular attraction

• Effect of temperature on solubility
• Generally, the solubility of solid solute in liquid solvents increases
with increase in temperature.
• It is obvious from graph that the
solubility increases with respect
temperature.

• Effect of temperature on solubility
There is no simple relationship between the
structure of a substance and the temperature
dependence of its solubility.

• Effect of temperature on solubility – other examples…

• Effect of temperature on solubility of gases
• The solubility of gases in liquids decreases with increasing
temperature.
• This is a reverse of the case of solids.
• Attractive intermolecular interactions in the gas phase are
essentially zero for most substances.

• When a gas dissolves, it does so because its molecules interact
with solvent molecules.
• Because heat is released when these new attractive interactions
form.
• Dissolving most gases in liquids is an exothermic process.
• Conversely, adding heat to the solution provides thermal energy
that overcomes the attractive forces between the gas and the
solvent molecules, thereby decreasing the solubility of the gas.

• The carbonated soft drinks have more bubbles
when they are in refrigerator.
• As temperature increases, gas molecules gain
more kinetic energy and are able to leave a
solution.

Cold

• Effect of pressure solubility
• The solubility of liquids and solids does not change appreciably with
pressure.
• But a solubility of gas in a liquid is directly proportional
• to its pressure.
• The concentration of gas molecules increases with
increasing pressure.
• The concentration of dissolved gas molecules in the
solution at equilibrium is also higher at higher pressures.

• The figure shows why the solubility increases with respect to pressure of
gas.

• (a) When a gas comes in contact with a pure liquid, some of the gas
molecules (purple spheres) collide with the surface of the liquid and
dissolve.
• When the concentration of dissolved gas molecules has increased so that
the rate at which gas molecules escape into the gas phase is the same as
the rate at which they dissolve.
• A dynamic equilibrium has been established, as depicted in figure.

• This equilibrium is entirely similar to the one that maintains the vapor
pressure of a liquid.
• (b) Increasing the pressure of the gas increases the number of molecules
of gas per unit volume, which increases the rate at which gas molecules
collide with the surface of the liquid and dissolve.
• (c) As additional gas molecules dissolve at the higher pressure, the
concentration of dissolved gas increases until a new dynamic equilibrium
is established.

• The relationship between pressure and the solubility of a gas is described
quantitatively by Henry’s law, which is named for its discoverer, the
English (British) physician and chemist, William Henry (1803).

• Effect of pressure solubility - Henry’s law
• Henry's law is one of the gas laws and was formulated by the British
physician / chemist, William Henry, in 1803.
“The solubility of a gas in a liquid is directly proportional
to the pressure of the gas”

Henry’s Law
• Henry's law has two major assumptions.
• Two major assumptions:
1. For the system is at low pressure.
2. For the species present as a very dilute solutions in liquid phase – low
concentration.

Henry’s Law – Other statements
• The partial pressure of the species in the vapour phase is directly
proportional to its liquid-phase mole fraction.
• At a constant temperature, the amount of a given gas dissolved in a
given type and volume of liquid is directly proportional to the partial
pressure of that gas in equilibrium with that liquid.
( )NiHxPy iii ,...,2,1==
pressureTotal:
constantsHenry':
fractionmolephase:
fractionmolephase:
P
H
Vy
Lx
i
i
i
−
−

Henry’s Law – continue
• Moreover, according to the Henry's law, the pressure of the gas is also
proportional to the solubility of gas in a particular solution / liquid
mixture at any constant temperature.
• The concentration of a solute gas in a solution is directly proportional to
the partial pressure of the gas over the solution. This can also be
represented as:
P = KHC
• Henry’s law is used to calculate the concentration of a gas in the
solution under pressure.
P is the partial pressure of the gas above the solution
KH is the Henry’s law constant for the solution
C is the concentration of the dissolved gas in solution

Henry’s Law constant
• Henry's law constant is an expression for equilibrium distribution of a
compound between liquid (water) and gaseous phase.
• Denoted by KH.
• Expressed as: KH = P / C
KH = atm / (mol / lit)
C = mol / lit
P = atm
• Different for every gas, temperature and solvent.

Henry’s Law – Validity and Limitations
• Henry's law is only an approximation that is applicable for dilute
solutions.
• Henry’s law is obeyed fairly satisfactorily by many gases of low
solubility, provided the pressure is not too high or the temperature is not
too low.
• The dissolved gas always follows the Henry’s law.
• The law is strictly applicable to those gases where the molecular species
are the same in the liquid phase.

Henry’s Law – Validity and Limitations
• It also applies simply for solutions where the solvent does not react
chemically with the gas being dissolved.
• Suggests that, Henry's law is related with physical absorption only.
• The gas phase should not undergo any chemical change.
• Large deviations from Henry’s law are observed in case of gases of
high solubility and particularly those which interact with the liquid
solvent.

Henry’s Law – Applications
• Soft drink industries
• Natural gas processing industries
• Various gas – liquid processes in chemical industries
• Miscellaneous

Henry’s Law – example problem
How many grams / quantity of carbon dioxide gas is dissolved in a 1
litre bottle of carbonated water if the manufacturer uses a pressure of
2.4 atm in the bottling process at 25 °C?
Given data:
KH of CO2 in water = 29.76 atm/(mol/L) at 25 °C
P = 2.4 atm;
We know the equation:
P = KHC
Therefore :
C = P / KH
Put these values…
2.4 / 29.76 = > 0.0806 mol / lit of CO2
Convert it into grams…

Henry’s Law – example problem
• 0.0806 mol / lit of CO2 (convert it into grams)
• 1 mole of CO2 = 12 + (16 × 2) = 12 + 32 = 44 g
• Mass of CO2 in grams = moles of CO2 × 44 g / mol
• Therefore :
= 0.0806 mol × 44 g / mol
= 3.546 g
There are 3.546 g of CO2 dissolved in a 1 litre bottle of carbonated water
from the manufacturer.
Molecular weight
of CO2

Henry’s Law – exercise problems…for you
To understand why soft drinks “fizz” and then go “flat” after being opened,
calculate the concentration of dissolved CO2 in a soft drink. The soft drink is
bottled under a pressure of 5.0 atm of CO2. In equilibrium with the normal
partial pressure of CO2 in the atmosphere. The Henry’s law constant for CO2
in water at 25°C is 3.4×10−2 M/atm.

• Effect of molecular size on solubility
• The larger the molecules of solute, the larger will be their molecular
weight and their size.
• It is more difficult for solvent molecules to surround the bigger molecules.
• Therefore, the larger molecules / particles are generally less soluble.

• Effect of stirring on solubility
• Stirring only increase the speed of the process.
• It increases the movement of the solvent that exposes the solute, thus
enabling the solubility.
• As molecules in the liquid substances are in constant move, the process
would take place any way, but it would take more time.

• Effect of intermolecular attraction on solubility
• The stronger the intermolecular attraction between solute and solvent, the
more likely the solute will dissolve.
• Example: ethanol in water, glucose in water – more soluble.
• Cyclohexane: not a water soluble.

• The study of various solutions / fluids / mixtures used in the chemical
process industries is very essential.
• The solution properties especially, physical / physicochemical /
thermophysical / thermodynamic properties of the solutions / mixtures
are very necessary to study.
• These properties can be of pure substances and the mixtures of various
substances.
• These properties are very important for the design, operation and
optimization of chemical processes.
• Chemistry of solutions – the solution properties

• There are various properties of the solutions as below:
• Density
• Viscosity
• Surface tension
• Refractive index
• Conductivity
• Resistivity
• Vapour pressure
Fundamental design
properties
• Chemistry of solutions – the solution properties…

• Most of the materials in a real world are not pure substances, but a mixture
of various types of substances.
• The pure substances from which solutions are prepared, are called as
components / constituents of the solutions.
• It is very important to study the properties of the pure components and
mixture of various substances.
• Since the temperature, pressure and compositions affects the solution
properties, therefore, it is essential to study the effect of these
parameters.

• Partial properties
The property of component when it is in mixture of one or more other
components.
• Derived properties
The property of the pure or mixture of solutions derived from another
basic or fundamental properties.
• Excess properties
The difference between actual property value of the solution and the
property value of the pure component.

organic and inorganic chemistry

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