Grade XI - First Lesson - Concepts in Chemistry

1. Basic concepts of Chemistry
Prepared by
Gowri Baskar

2. Roald Hoffmann
• Chemistry is the science of molecules and
their transformations. It is the science not so
much of the one hundred elements but of the
infinite variety of molecules that may be built
from them.

3. The branch of science that studies the
preparation, properties, structure and
reactions of material substances is called
chemistry.

4. Two interesting things
• Philosopher’s stone
• Elixir of life

5. Philosophers stone
• Is a legendary alchemical substance capable of
turning base metals such as mercury into gold.
• The philosophers' stone is created by the
alchemical method known as The Magnum
Opus or The Great Work.

6. Elixir of life
• Also known as elixir of immortality
• Is a potion that supposedly grants the drinker
eternal life and/or eternal youth.
• Said to cure all diseases.
• Alchemists in various ages and cultures sought
the means of formulating the elixir.
• The concept originated in ancient India or
China where the concept preceded that in
Europe by millennia.

7. Chemistry in India
• Other names - Rasayan Shastra, Rastantra,
Ras Kriya or Rasvidya.
• Includes metallurgy, medicine, manufacture of
cosmetics, glass, dyes
• Examples - Systematic excavations at
Mohenjodaro in Sindh and Harappa in Punjab

8. Advancements in olden days
• Construction work – baked bricks, mass
production of pottery, gypsum cement
• Falence – sort of glass used in ornaments
• Improved hardness of copper
• Copper metallurgy in India dates back to the
beginning of chalcolithic cultures in the
subcontinent. There are much archeological
evidences to support the view that
technologies for extraction of copper and iron
were developed indigenously.

9. • Rig veda – Tanning of leather
• Golden gloss of black polished ware of
northern India
• Kautilya’s Arthasasthra – Production of salt
from sea
• Charaka Samhita - how to prepare sulphuric
acid, nitric acid and oxides of copper, tin and
zinc; the sulphates of copper, zinc and iron
and the carbonates of lead and iron

10. Modern Chemistry
• Took shape in the 18th century Europe, after a
few centuries of alchemical traditions which
were introduced in Europe by the Arabs.
• Especially the Chinese and the Indian – had
their own alchemical traditions. These
included much knowledge of chemical
processes and techniques.

11. Father of chemistry
• Antoine Laurent Lavoisier
• French nobleman and
chemist
• 18th century – Chemical
Revolution
• Has influence in both
chemistry and biology

12. Chemistry
Inorganic
Organic
Physical
Analytical
Biochemistry

13. • Organic Chemistry: Study of carbon and its
compounds
• Inorganic Chemistry: Study of compounds not
covered by organic chemistry or compounds that
don't contain a C-H bond
• Analytical Chemistry: Study of the chemistry of
matter and the development of tools to measure
properties of matter
• Physical Chemistry: Branch of chemistry that
applies physics to the study of chemistry - the
applications of thermodynamics and quantum
mechanics to chemistry
• Biochemistry: Study of chemical processes that
occur inside of living organisms

14. Matter
• Anything that possess mass, occupies space,
offers resistance and can be perceived by one
or more of our senses is called Matter.
• Matter is made up of particles.
• Particles will have space between them and
continuously moving and attract each other.

15. Matter
Physical
classification
Solid Liquid Gas
Chemical
classification
Pure
substances
Elements
Metals Non-Metals Metalloids
Compounds
Organic
compounds
Inorganic
compounds
Mixtures

16. Elements
• It is the simplest form of pure substance
• It can neither be decomposed into nor built
from simpler substances by ordinary physical
and chemical methods.
• It contains only one kind of atoms.
• The number of elements known till date is
118.
• An element can be a metal, a non-metal or a
metalloid.

17. Facts about metals
• Hydrogen is the most abundant element in the
universe.
• OXYgen (46.6%), a non-metal, is the most
abundant element in the earth crust.
• AI is the most abundant metal in the earth crust.

18. Symbol
• A symbol is an abbreviation or
shortened form for the full name
of an element.
• The present system of symbols
was introduced by Berzelius.
• Baron Jöns Jacob Berzelius was a
Swedish chemist.

19. Element Symbol Name
Sodium Na Natrium
Potassium K Kalium
Antimony Sb Stibium
Copper Cu Cuprum
Gold Au Aurum
Silver Ag Argentum
Iron Fe Ferum
Lead Pb Plumbum
Mercury Hg Hydragyrum
Tin Sn Stannum
Tungsten W Wolfram

20. Compounds
• It is a form of matter which can be formed by
combining two or more elements in a definite
ratio by mass.
• It can be decomposed into its constituent
elements by suitable chemical methods, e.g.,
water (H2O) is made of hydrogen and oxygen
in the ratio 1 : 8 by mass.

21. Inorganic compounds
• Previously, it was believed
that these compounds are
derived from non-living
sources, like rocks and
minerals.
• But these are infact the
compounds of all the
elements except hydrides of
carbon (hydrocarbons) and
their derivatives.
Organic compounds
• According to earlier
scientists, these compounds
are derived from living
sources like plants and
animals, or these remain
buried under the earth
(e.g., petroleum).
• According to modern
concept, these are the
hydrides of carbon and their
derivatives.

22. • A mixture contains particles of two or more
pure substances which may be present in it in
any ratio.
• In a homogeneous mixture, the components
completely mix with each other.
• In a heterogeneous mixture, the composition
is not uniform throughout and sometimes
different components are visible.

23. Common separation techniques of
mixtures
• Chromatography is the separation of a
mixture by passing it in solution or suspension
or as a vapor (as in gas chromatography)
through a medium in which the components
move at different rates.

24. Distillation
• It is an effective method to separate mixtures
comprised of two or more pure liquids.
Distillation is a purification process where the
components of a liquid mixture are vaporized
and then condensed and isolated.

25. • Evaporation is a technique used to separate
out homogenous mixtures where there is one
or more dissolved solids. This method drives
off the liquid components from the solid
components.

26. Filtration
• It is a separation method used to separate out
pure substances in mixtures comprised of
particles some of which are large enough in
size to be captured with a porous material.
Particle size can vary considerably, given the
type of mixture.

27. • Sublimation This is the process of conversion of a
solid directly into vapours on heating. Substances
showing this property are called sublimate, e.g.,
iodine, naphthalene, camphor. This method is
used to separate a sublimate from non-sublimate
substances.
• Crystallisation It is a process of separating solids
having different solubilities in a particular solvent.
• Magnetic separation Tills process is based upon
the fact that a magnet attracts magnetic
components of a mixture of magnetic and non-
magnetic substances.

28. Atoms and Molecules

29. • An atom
– Is the smallest particle of the element that can
exist independently and retain all of its chemical
properties.
– It is made up of fundamental particles like
electrons, protons and neutrons.
• A Molecule
– It is generally a group of one or more atoms that
are chemically bonded and are tightly held
together.
– Smallest particle of an element or compound.
– Constituted by same type of atoms.

31. Chemical and Physical properties

32. • We need to measure all physical quantities.
We can express the value of a physical
quantity as the product of the numerical value
and the unit in which it is expressed.
• Fundamental units:
– Which can neither be derived from one another
nor they can be further resolved into any other
units.

33. Fundamental Units
Quantity Name of Unit Abbreviation
Mass Kilogram Kg
Length Meter M
Temperature Kelvin or Celsius K or C
Amount of
substance
Mole Mol
Time Second S
Electric current Ampere A
Luminous
Intensity
Candela Cd

34. Derived Units
Quantity S.I unit Symbol
Velocity Meter per second m/s
Area Square meter m2
Volume Cubic meter m3
Density Kilogram/m3 Kg/m3
Energy Joule Kgm2/s2
Force Newton Kgm/s2
Electrical charge Coulomb Ampere - second

35. Definition – Unit of length - metre
• The metre is the length of the path travelled
by light in vacuum during a time interval of
1/299 792 458 of a second.
Definition – Unit of mass - kilogram
• The kilogram is the unit of mass; it is equal
to the mass of the international prototype
of the kilogram.

36. Definition – Unit of time - second
• The second is 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 caesium-133 atom.
Definition – Unit of electric current -
Ampere
The ampere is that constant current, which if maintained
in two straight parallel conductors of infinite length of
negligible circular cross-section and placed 1 metre apart
in vacuum, would produce between these conductors a
force equal to 2 × 10–7 newton per metre of length.

37. Definition – Unit of thermodynamic
temperature - kelvin
• The kelvin, unit of thermodynamic temperature,
is the fraction 1/273.16 of the thermodynamic
temperature of the triple point of water.
Definition – Unit of Luminous
intensity - Candela
The candela is the luminous intensity, in a given direction,
of a source that emits monochromatic radiation of
frequency 540 × 1012 hertz and that has a radiant intensity
in that direction of 1/683 watt per steradian.

38. Definition – Unit of amount of
substance - mole
• 1. The mole is the amount of substance of a
system, which contains as many elementary
entities as there are atoms in 0.012 kilogram
of carbon-12; its symbol is ‘mol’.
• 2. When the mole is used, the elementary
entities must be specified and these may be
atoms, molecules, ions, electrons, other
particles, or specified groups of such particles.

40. Mass and weight
Mass Weight
Amount of matter present in a body Force exerted by gravity on a object
Scalar quantity Vector quantity
SI unit – kg SI unit – N
Constant regardless of gravitational force Varies with gravitational field strength
Measured by beam balance or calibrated
electronic balance
Measured by spring or compression
balance
Cannot be zero Can be zero if no gravity

41. Volume
• Volume is the amount of space occupied by a
substance
• It has the units of (length)3.
• A common unit, litre (L) which is not an SI
unit, is used for measurement of volume of
liquids.

42. Different units to express volume

43. Volume measuring devices
• A volumetric flask is used to prepare a known
volume of a solution.

44. Density
• Density of a substance is its amount of mass
per unit volume.
• SI unit of density – kg/m3 or g/cm3
• Density of a substance tells us about how
closely its particles are packed.
• If density is more, it means particles are more
closely packed.

45. Temperature
• There are three common scales to measure
temperature
– °C (degree celsius)
– °F (degree fahrenheit)
– K (kelvin).
• Here, K is the SI unit.

46. • The temperatures on two scales are related to
each other by the following relationship:
• It is interesting to note that temperature
below 0 °C (i.e., negative values) are possible
in Celsius scale but in Kelvin scale, negative
temperature is not possible.

47. Platinum- Iridium alloy
• It is used as a standard for both mass and length.
• It is the most inert metal.
• It doesn't react with any other
compound/atmospheric oxygen forming oxide.
• So the weight of the metal stays constant, that is it
doesn't increase or decrease, providing a
“standard” weight.
• Also, platinum is malleable and ductile.

48. Uncertainty in measurement

49. Key Terms
• Scientific notation: A way of writing numbers
that are too big or too small to be
conveniently written in standard form.
• Integer: An element of the infinite and
numerable set {…,-3,-2,-1,0,1,2,3,…}.
• Order of Magnitude: An order of magnitude is
the class of scale or magnitude of any amount,
where each class contains values of a fixed
ratio to the class preceding it.

50. Scientific notation
• To express a number in scientific notation, you
move the decimal place to the right if the number
is less than zero or to the left if the number is
greater than zero.
• The decimal would move five places to the left to
get 456000 as 4.56×105
• Similarly, 0.00016 can be written as 1.6 × 10–4.
Here, the decimal has to be moved four places to
the right and (–4) is the exponent in the scientific
notation..

51. Mathematical operations
• Multiplication and Division

52. • Addition and Subtraction
– Numbers must be written in same pattern or way

53. Significant figures
• Every experimental measurement has some
amount of uncertainty associated with it
because of limitation of measuring instrument
and the skill of the person making the
measurement.
• For example, mass of an object is obtained
using a platform balance and it comes out to
be 9.4g. On measuring the mass of this object
on an analytical balance, the mass obtained is
9.4213g.

54. • The uncertainty in the experimental or the
calculated values is indicated by mentioning
the number of significant figures.
• Significant figures are meaningful digits which
are known with certainty plus one which is
estimated or uncertain.
• The uncertainty is indicated by writing the
certain digits and the last uncertain digit.

55. Certain rules for determining the
number of significant figures
• All non-zero digits are significant. For example
in 285 cm, there are three significant figures
and in 0.25 mL, there are two significant
figures.
• Zeros preceding to first non-zero digit are not
significant. Such zero indicates the position of
decimal point. Thus, 0.03 has one significant
figure and 0.0052 has two significant figures.
• Zeros between two non-zero digits are
significant. Thus, 2.005 has four significant
figures.

56. • Zeros at the end or right of a number are
significant, provided they are on the right side
of the decimal point.
• Counting the numbers of object, for example,
2 balls or 20 eggs, have infinite significant
figures as these are exact numbers and can be
represented by writing infinite number of
zeros after placing a decimal i.e., 2 = 2.000000
or 20 = 20.000000.

57. Precision and Accuracy
• Precision refers to the closeness of various
measurements for the same quantity.
• Accuracy is the agreement of a particular
value to the true value of the result.

58. Rounding off results
• If the rightmost digit to be removed is more than 5, the
preceding number is increased by one. For example,
1.386. If we have to remove 6, we have to round it to
1.39.
• If the rightmost digit to be removed is less than 5, the
preceding number is not changed. For example, 4.334
if 4 is to be removed, then the result is rounded upto
4.33.
• 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 6.35 is to be rounded by
removing 5, we have to increase 3 to 4 giving 6.4 as the
result. However, if 6.25 is to be rounded off it is
rounded off to 6.2.

59. Dimensional Analysis
• There is a need to convert units from one
system to the other. The method used to
accomplish this is called factor label method
or unit factor method or dimensional analysis.
• Ex: A piece of metal is 3 inch (represented by
in) long. What is its length in cm?

60. LAWS OF CHEMICAL COMBINATIONS
• The laws of chemical combination describe
the basic principles obeyed by interacting
atoms and molecules, interactions that can
include many different combinations that
happen in many different ways.
• This amazing diversity of interactions allows
for an astounding variety of chemical
reactions and compounds.

61. Law of Conservation of Mass
• Was formulated by Antoine Lavoisier in 1789.
• “The law of conservation of mass states that
the net change in mass of the reactants and
products before and after a chemical reaction
is zero. This means mass can neither be
created nor destroyed. In other words, the
total mass in a chemical reaction remains
constant.”

63. Law of Definite/Constant proportions
• Was formulated by Joseph Proust.
• “He stated that a given compound always
contains exactly the same proportion of
elements by weight.”

64. • Proust worked with two samples of cupric
carbonate — one of which was of natural
origin and the other was synthetic.
• Thus, he concluded that irrespective of the
source, a given compound always contains
same elements combined together in the
same proportion by mass.

65. Law of multiple proportions
• Was proposed by John Dalton in 1803.
• According to this law, “ if two elements can
combine to form more than one compound,
the masses of one element that combine with
a fixed mass of the other element, are in the
ratio of small whole numbers.”

67. Gay Lussac’s law of Gaseous volumes
• Was given by Gay Lussac in 1808.
• “When gases combine or are produced in a
chemical reaction they do so in a simple ratio
by volume, provided all gases are at the same
temperature and pressure.”

68. • 100 mL of hydrogen combine with 50 mL of
oxygen to give 100 mL of water vapour.
• Thus, the volumes of hydrogen and oxygen
which combine (i.e., 100 mL and 50 mL) bear a
simple ratio of 2:1.
• The Gay Lussac’s law was explained properly
by the work of Avogadro in 1811.

69. Avogadro’s Law
• In 1811, Avogadro proposed that equal
volumes of all gases at the same temperature
and pressure should contain equal number of
molecules.
• Avogadro made a distinction between atoms
and molecules which is quite understandable
in present times.

71. Dalton’s Atomic Theory
• The origin of the idea that matter is
composed of small indivisible particles called
‘atomio’ (meaning, indivisible), dates back to
the time of Democritus, a Greek Philosopher
(460–370 BC),

72. • In 1808, Dalton published ‘A New System of
Chemical Philosophy’, in which he proposed
the following :
– Matter consists of indivisible atoms.
– All atoms of a given element have identical
properties, including identical mass. Atoms of
different elements differ in mass.
– Compounds are formed when atoms of different
elements combine in a fixed ratio.
– Chemical reactions involve reorganisation of
atoms. These are neither created nor destroyed in
a chemical reaction.