Some Basic Concepts of Chemistry

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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. - Roald Hoffmann

Roald Hoffmann (born Roald Safran; July 18, 1937) is a Polish-American theoretical chemist who won the 1981 Nobel Prize in Chemistry.

For the sake of convenience, science is sub-divided into various disciplines: chemistry, physics, biology, geology, etc.

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

Development of chemistry Chemistry, as we understand it today, is not a very old discipline. Chemistry was not studied for its own sake, rather it came up as a result of search for two interesting things:

i. Philosopher’s stone (Paras) which would convert all baser metals e.g., iron and copper into gold.
ii. ‘Elexir of life’ which would grant immortality.
People in ancient India, already had the knowledge of many scientific phenomenon much before the advent of modern science. They applied that knowledge in various walks of life.

Chemistry developed mainly in the form of Alchemy and Iatrochemistry during 1300-1600 CE.

What is alchemy in chemistry? Alchemy is an old study of changing basic substances (such as metals) to other substances. “The word alchemy comes from the Arabic al-kimiyya, which in turn probably comes from Greek.”

Modern chemistry took shape in the 18th century Europe, after a few centuries of alchemical traditions which were introduced in Europe by the Arabs.
In ancient India, chemistry was called Rasayan Shastra, Rastantra, Ras Kriya or Rasvidya.
It included metallurgy, medicine, manufacture of cosmetics, glass, dyes, etc. Systematic excavations at Mohenjodaro in Sindh and Harappa in Punjab prove that the story of development of chemistry in India is very old. Archaeological findings show that baked bricks were used in construction work. It shows the mass production of pottery, which can be regarded as the earliest chemical process, in which materials were mixed, moulded and subjected to heat by using fire to achieve desirable qualities.
Remains of glazed pottery have been found in Mohenjodaro.
Gypsum cement has been used in the construction work. It contains lime, sand and traces of CaCO₃.
Harappans made faience, a sort of glass which was used in ornaments. They melted and forged a variety of objects from metals, such as lead (Pb), silver(Ag), gold(Au) and copper(Cu).
They improved the hardness of copper for making artefacts by using tin (Sn) and arsenic (As).
A number of glass objects were found in Maski (Karnataka) in South India (1000–900 BCE), and Hastinapur and Taxila in North India (1000–200 BCE).
Glass and glazes were coloured by addition of colouring agents like metal oxides.
Copper metallurgy in India dates back to the beginning of chalcolithic cultures in the subcontinent.

What is metallurgy and its uses? Metallurgy is defined as a process that is used for the extraction of metals in their pure form. The compounds of metals mixed with soil, limestone, sand, and rocks are known as minerals. Metals are commercially extracted from minerals at low cost and minimum effort.

There are much archeological evidences to support the view that technologies for extraction of copper and iron were developed indigenously.

Archaeology is the study of the societies and peoples of the past by examining the remains of their buildings, tools, and other objects.

According to Rigveda, tanning of leather and dying of cotton were practised during 1000–400 BCE. The golden gloss of the black polished ware of northen India could not be replicated and is still a chemical mystery.

What does term gloss mean? a surface luster or brightness

These wares indicate the mastery with which kiln temperatures could be controlled.

What is the meaning of the word kiln?
A furnace or oven for drying, burning, or baking something, as bricks, grain, or pottery.

Kautilya’s Arthashastra describes the production of salt from sea.
A vast number of statements and material described in the ancient Vedic literature can be shown to agree with modern scientific findings.
Copper utensils, iron, gold, silver ornaments and terracotta discs and painted grey pottery have been found in many archaeological sites in north India.

Terracotta - a type of fired clay, typically of a brownish-red colour and unglazed, used as an ornamental building material and in modelling.

Sushruta Samhita explains the importance of Alkalies.
The Charaka Samhita mentions ancient indians who knew 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.

Rasopanishada describes the preparation of gunpowder mixture. Tamil texts also describe the preparation of fireworks using sulphur, charcoal, saltpetre (i.e., potassium nitrate), mercury, camphor, etc.
Nagarjuna was a great Indian scientist. He was a reputed chemist, an alchemist and a metallurgist.
His work Rasratnakar deals with the formulation of mercury compounds.

'Rasaratnakar' and 'Arogyamanjari' were books written by Acharya Nagarjun on and respectively.

He has also discussed methods for the extraction of metals, like gold, silver, tin and copper. A book, Rsarnavam, appeared around 800 CE.

Rsarnavam discusses the uses of various furnaces, ovens and crucibles for different purposes. It describes methods by which metals could be identified by flame colour.
The credit for inventing soap also goes to him. He used mustard oil and some alkalies as ingredients for making soap. Indians began making soaps in the 18th century CE. Oil of Eranda and seeds of Mahua plant and calcium carbonate were used for making soap. The paintings found on the walls of Ajanta and Ellora, which look fresh even after ages, testify to a high level of science achieved in ancient India.

Chakrapani discovered mercury sulphide.

Write formula of mercury sulphide. Mercury sulfide, or Mercury(II) sulfide is a chemical compound composed of the chemical elements mercury and sulfur. It is represented by the chemical formula HgS. Mercury(II) sulfide, HgS, is a black or red crystalline solid used chiefly as a pigment in paints, rubber, and plastics.
Mercurous chloride - Hg₂Cl₂

Varähmihir’s Brihat Samhita is a sort of encyclopaedia, which was composed in the sixth century CE. It informs about the preparation of glutinous material to be applied on walls and roofs of houses and temples.
Glutinous means glue-like, or thick and sticky.
It was prepared entirely from extracts of various plants, fruits, seeds and barks, which were concentrated by boiling, and then, treated with various resins. It will be interesting to test such materials scientifically and assess them for use.

A number of classical texts, like Atharvaveda (1000 BCE) mention some dye stuff, the material used were turmeric, madder, sunflower, orpiment, cochineal and lac. Some other substances having tinting property were kamplcica, pattanga and jatuka.

Varähmihir’s Brihat Samhita gives references to perfumes and cosmetics.

Recipes for hair dying were made from plants, like indigo and minerals like iron power, black iron or steel and acidic extracts of sour rice gruel.
Gandhayukli describes recipes for making scents, mouth perfumes, bath powders, incense and talcum power. Paper was known to India in the 17th century as account of Chinese traveller I-tsing describes.
Excavations at Taxila indicate that ink was used in India from the fourth century.
Colours of ink were made from chalk, red lead and minium.

Minium, also known as red lead or red lead oxide, is a bright orange red pigment

It seems that the process of fermentation was well-known to Indians. Vedas and Kautilya’s Arthashastra mention about many types of liquors. Charaka Samhita also mentions ingredients, such as barks of plants, stem, flowers, leaves, woods, cereals, fruits and sugarcane for making Asavas.

Asava or Asavam is a category of Ayurvedic medicine prepared by the fermentation of fresh herbal juices

The concept that matter is ultimately made of indivisible building blocks, appeared in India a few centuries BCE as a part of philosophical speculations.
Acharya Kanda, born in 600 BCE, originally known by the name Kashyap, was the first proponent of the ‘atomic theory’. He formulated the theory of very small indivisible particles, which he named ‘Paramãnu’ (comparable to atoms).

Acharya Kanda authored the text Vaiseshika Sutras.

According to him, all substances are aggregated form of smaller units called atoms (Paramãnu), which are eternal, indestructible, spherical, suprasensible and in motion in the original state.
He explained that this individual entity cannot be sensed through any human organ. Kanda added that there are varieties of atoms that are as different as the different classes of substances.
He said these (Paramãnu) could form pairs or triplets, among other combinations and unseen forces cause interaction between them.
He conceptualised this theory around 2500 years before John Dalton (1766-1844).

Charaka Samhita is the oldest Ayurvedic epic of India.

It describes the treatment of diseases. The concept of reduction of particle size of metals is clearly discussed in Charaka Samhita. Extreme reduction of particle size is termed as nanotechnology.

Nanotechnology, the manipulation and manufacture of materials and devices on the scale of atoms or small groups of atoms.
Charaka Samhita describes the use of bhasma of metals in the treatment of ailments. Now-a-days, it has been proved that bhasmas have nanoparticles of metals.

After the decline of alchemy, Iatrochemistry reached a steady state, but it too declined due to the introduction and practise of western medicinal system in the 20th century.

Iatrochemistry a school of thought of the 16th and 17th centuries which sought to understand medicine and physiology in terms of chemistry.

During this period of stagnation, pharmaceutical industry based on Ayurveda continued to exist, but it too declined gradually. It took about 100-150 years for Indians to learn and adopt new techniques.
During this time, foreign products poured in. As a result, indigenous traditional techniques gradually declined. Modern science appeared in Indian scene in the later part of the nineteenth century. By the mid-nineteenth century, European scientists started coming to India and modern chemistry started growing.

Importance of Chemistry

Why is Chemistry important to us?
Chemistry is the field of science that studies the characteristics, content, and structure of elements and compounds, as well as how they change and the energy generated or absorbed as a result of those changes.
Importance in society: 1. Principles of chemistry are applicable in diverse areas, such as weather patterns, functioning of brain and operation of a computer, production in chemical industries, manufacturing fertilisers, alkalis, acids, salts, dyes, polymers, drugs, soaps, detergents, metals, alloys, etc., including new material.
2. Chemistry contributes in a big way to the national economy.
a. large-scale production of a variety of fertilisers, improved variety of pesticides and insecticides.
b. isolation of lifesaving drugs from natural sources and makes possible synthesis of such drugs, e.g. cisplatin and taxol, which are effective in cancer therapy; AZT (Azidothymidine) is used for helping AIDS patients.
c. Chemistry contributes to a large extent in the development and growth of a nation. Various industries contribute in a big way to the economy of a nation and generate employment, e.g. the production of superconducting ceramics, conducting polymers, optical fibres, manufacture of utility goods like acids, alkalies, dyes, polymesr metals, etc.
d. Useful in pressing aspects of environmental degradation-
Safer alternatives to environmentally hazardous refrigerants, like CFCs (chlorofluorocarbons), responsible for ozone depletion in the stratosphere, have been successfully synthesised.
- However, many big environmental problems include the management of the Green House gases, like methane, carbon dioxide, etc.
- Understanding of biochemical processes, use of enzymes for large-scale production of chemicals etc.

8. Importance of Chemistry
Chemistry has a direct impact on our life and has wide range of applications in different fields. These are given below:
→ (a) In Agriculture and Food:
(i) It has provided chemical fertilizers such as urea, calcium phosphate, sodium nitrate, ammonium phosphate etc. It has helped to protect the crops from insects and harmful bacteria, by the use ‘of certain effective insecticides, fungicides and pesticides.
(ii) The use of preservatives has helped to preserve food products like jam, butter, squashes etc. for longer periods.
(b) In Health and Sanitation:
(i) Life-saving drugs. Cisplatin and taxol have been found to be very effective for cancer therapy and AZT (Azidothymidine) is used for AIDS victims.
(ii) Disinfectants such as phenol are used to kill the micro-organisms present in drains, toilet, floors etc.
(iii) A low concentration of chlorine i.e., 0.2 to 0.4 parts per million (ppm) is used ’ for sterilization of water to make it fit for drinking purposes.
(c) Saving the Environment:
The rapid industrialisation all over the world has resulted in lot of pollution. Poisonous gases and chemicals are being constantly released in the atmosphere. They are polluting environment at an alarming rate. Scientists are working day and night to develop substitutes which may cause lower pollution. For example, CNG (Compressed Natural Gas), a substitute of petrol, is very effective in checking pollution caused by automobiles.
(d) Application in Industry:
Chemistry has played an important role in developing many industrially manufactured fertilizers, alkalis, acids, salts, dyes, polymers, drugs, soaps, detergents, metal alloys and other inorganic and organic chemicals including new materials contribute in a big way to the national economy. ↔

1. In chemical classification, the matter is classified as elements, mixtures, and compounds.
2. A homogeneous mixture has uniform composition and properties throughout while a heterogeneous mixture’s composition and properties are not uniform throughout.

→ Matter
Anything which has mass and occupies space is called matter.
For example, book, pencil, water, air are composed of matter as we know that they have mass and they occupy space.
Classification of Matter
(A) Physical classification
(B) Chemical classification
(A) Physical Classification:
Matter can exist in three physical states:
1. Solids 2. Liquids 3. Gases
1. Solids: The particles are held very close to each other in an orderly fashion and there is not much freedom of movement.
Characteristics of solids: Solids have definite volume and definite shape.
2. Liquids: In liquids, the particles are close to each other but can move around.
Characteristics of liquids: Liquids have definite volume but not definite shape.
3. Gases: In gases, the particles are far apart as compared to those present in solid or liquid states. Their movement is easy and fast.
Characteristics of Gases: Gases have neither definite volume nor definite shape. They completely occupy the container in which they are placed. ↔
(B) Chemical Classification:
Based upon the composition, matter can be divided into two main types:
→ 1. Pure Substances 2. Mixtures.
1. Pure substances: A pure substance may be defined as a single substance (or matter) which cannot be separated by simple physical methods.
Pure substances can be further classified as (i) Elements (ii) Compounds
(i) Elements: An element consists of only one type of particles. These particles may be atoms or molecules.
e.g. sodium, copper, silver, hydrogen, oxygen etc. are some examples of elements. ↔
They all contain atoms of one type. However, atoms of different elements are different in nature. Some elements such as sodium or copper contain single atoms held together as their constituent particles whereas in some others two or more atoms combine to give molecules of the element. Thus, hydrogen, nitrogen and oxygen gases consist of molecules in which two atoms combine to give the respective molecules of the element.

→ (ii) Compounds:
It may be defined as a pure substance containing two or more elements combined together in a fixed proportion by weight and can be decomposed into these elements by suitable chemical methods. Moreover, the properties of a compound are altogether different from the constituting elements.
Two types
(i) Inorganic Compounds: These are compounds which are obtained from non-living sources such as rocks and minerals. A few examples are: Common salt, marble, gypsum, washing soda etc.
(ii) Organic Compounds are the compounds which are present in plants and animals. All the organic compounds have been found to contain carbon as their essential constituent. For example, carbohydrates, proteins, oils, fats etc. ↔
→ 2. Mixtures: The combination of two or more elements or compounds which are not chemically combined together and may also be present in any proportion, is called mixture. A few examples of mixtures are: milk, sea water, petrol, lime water, paint glass, cement, wood etc.
Types of mixtures: Mixtures are of two types:
(i) Homogeneous mixtures: A mixture is said to be homogeneous if it has a uniform composition throughout and there are no visible boundaries of separation between the constituents.
e.g. A mixture of sugar solution in water has the same sugar water composition throughout and all portions have the same sweetness.
(ii) Heterogeneous mixtures: A mixture is said to be heterogeneous if it does not have uniform composition throughout and has visible boundaries of separation between the various constituents. The different constituents of a heterogeneous mixture can be seen even with naked eye.
e.g. When iron filings and sulphur powder are mixed together, the mixture formed is heterogeneous. It has greyish-yellow appearance and the two constituents, iron and sulphur, can be easily identified with naked eye. ↔
→ Differences between Compounds and Mixtures
1. In a compound, two or more elements are combined chemically. In a mixture, or more elements or compounds are simply mixed and not combined chemically.
2. In a compound, the elements are present in the fixed ratio by mass. This ratio cannot change. In a mixture the constituents are not present in fixed ratio. It can vary.
3. Compounds are always homogeneous i.e., they have the same composition throughout. Mixtures may be either homogeneous or heterogeneous in nature.
4 In a compound, constituents cannot be separated by physical methods. Constituents of mixtures can be separated by physical methods.
5. In a compound, the constituents lose their identities i.e., compound does not show the characteristics of the constituting elements. In a mixture, the constituents do not lose their identities i.e., a mixture shows the characteristics of all the constituents. ↔

→ Properties of Matter and Their Measurements
Physical Properties:
Those properties which can be measured or observed without changing the identity or the composition of the substance.
Some examples of physical properties are colour, odour, melting point, boiling point etc.
Chemical Properties: It requires a chemical change to occur. The examples of chemical properties are characteristic reactions of different substances. These include acidity, basicity, combustibility etc. ↔

→ 5392 Click Here-Seven basic units; Pre fix in SI Units

Units of Measurement

Fundamental Units:
The quantities mass, length and time are called fundamental quantities and their units are known as fundamental units.
Si-System: This system of measurement is the most common system employed throughout the world.
It has given units of all the seven basic quantities listed above. The International System of Units is a measuring system based on seven basic units
length, mass, time, temperature, amount of substance, electric current and luminous intensity.
The International System of Units (SI), sometimes referred to as the metric system, is the international metric standard used for measuring. On May 20, in the year of 1875, seventeen nations, along with the United States, signed the International Treaty of the Meter in Paris. This day is recognized across the world as World Metrology Day. ↔
The National Institute of Standards and Technology (NIST) represents the United States in the numerous international agencies set in place by the Meter General meeting.

Centimeter-Gram-Seconds: The centimeter–gram–second unit system is a version of the system of measurement program that utilizes the centimeter as a measurement of distance length, the gram as a measuring unit of mass, and the second as a unit of time.
Foot-pound-second (fps system of units): The foot-pound-second (fps) system of units is a scheme for measuring dimensional and material quantities. The fundamental units are the foot for length, the pound for weight, and the second for time.

Meter-Kilogramme-Second: Meter-Kilogramme-Second (MKS) is a measurable unit system based on the meter, kilogram, and second as quantities of length, mass, and time; it is the foundation of the SI units. ↔

Definitions of Basic SI Units
1. Metre: It is the length of the path travelled by light in vacuum during a time interval of 1/299792458 of a second.
2. Kilogram: It is the unit of mass. It is equal to the mass of the international prototype of the kilogram.
3. Second: It is the duration of 9192631, 770 periods of radiation which correspond to the transition between the two hyper fine levels of the ground state of caesium- 133 atom.
4. Kelvin: It is the unit of thermodynamic temperature and is equal to 1/273.16 of the thermodynamic temperature of the triple point of water.
5. 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 x 10^-7 N per metre of length.
6. Candela: It may be defined as the luminous intensity in a given direction, from a source which emits monochromatic radiation of frequency 540 x 10¹² Hz and that has a radiant intensity in that direction of 1/ 683 watt per steradian.
7. Mole: It is the amount of substance which contains as many elementary entities as there are atoms in 0.012 kilogram of carbon -12. Its symbol is ‘mol’.
Note- Triple point of water is 0.01 °C or 279.16K (32.01°F).

Mass and Weight
Mass: Mass of a substance is the amount of matter present in it. The mass of a substance is constant.
The mass of a substance can be determined accurately in the laboratory by using an analytical balance.
SI unit of mass is kilogram.
Weight: It is the force exerted by gravity on an object. Weight of substance may vary from one place to another due to change in gravity.
Volume: Volume means the space occupied by matter. It has the units of (length)3. In SI units, volume is expressed in metre³ (m³). However, a popular unit of measuring volume, particularly in liquids is litre (L) but it is not in SI units or an S.I. unit. Mathematically,
→ 1L = 1000 mL = 1000 cm³ = 1dm³. ↔

Analytical Balance-

What is the difference between physical and chemical balance?

Chemical balance is more sensitive than physical balance. It can weigh up to 4 places of decimals.

What is a rider? What is its weight?

Riders is used to determine the masses of objects or samples, or in other words the quantity of things in them. Equilibrium is a tool of comparison. It compares normal or known weights on the left panel to unknown specimens or items on the right panel. A rider weighs 10mg.

Why is substances not directly weighed on the chemical balance?

An analytical balance is so vulnerable that a single grain mass of a chemical substance can be detected. Therefore, if a direct weighing technique is used, the material should be added to the tared container that holds it, NEVER straight to the pan or even weighing paper on the pan.

Why is the front door of the balance closed during weighing?

Close the balancing door while weighing an item to avoid disturbing reading by air currents. In order to avoid dust and dirt from entering the equilibrium, the operator should close the balancing gate when completed.

What is the standard solution?

A standard solution in analytical chemistry is a solution that contains an accurately recognized concentration of an element or a substance. To create a particular quantity, a known solvent weight is dissolved. It is prepared using a standard substance, such as a primary standard.
Volume of liquids can be measured by different devices like burette, pipette, graduated cylinder, volumetric flask etc. All of them have been calibrated.

Burette is laboratory apparatus used in quantitative chemical analysis to measure the volume of a liquid or a gas. It consists of a graduated glass tube with a stopcock (turning plug, or spigot) at one end.
A pipette (sometimes spelled as pipet) is a laboratory tool commonly used in chemistry, biology and medicine to transport a measured volume of liquid, often as a media dispenser.
A graduated cylinder, also known as a measuring cylinder or mixing cylinder, is a common piece of laboratory equipment used to measure the volume of a liquid. It has a narrow cylindrical shape. Each marked line on the graduated cylinder represents the amount of liquid that has been measured.
A volumetric flask (measuring flask or graduated flask) is a piece of laboratory apparatus, a type of laboratory flask, calibrated to contain a precise volume at a certain temperature. Volumetric flasks are used for precise dilutions and preparation of standard solutions.

Density
Density is the measurement of how tightly a material is packed together. It is defined as the mass per unit volume.
Density = Mass / Volume Density of a substance is its amount of mass per unit volume. However, the SI unit of Density is measured using kilograms per cubic metre (kg/m₃).
A few other units are
gram per millilitre (g/mL)
kilogram per litre (kg/L) gram per cubic centimetre (g/cm₃)
1 g/cm₃ = 1000 kg/m₃ kilogram per cubic decimetre (kg/dm₃) In the cgs system density is measured in grams per cubic centimetre (g/cm₃).
→ Temperature:
There are three scales in which temperature can be measured. These are known as Celsius scale (°C), Fahrenheit scale (°F) and Kelvin scale (K). ↔
The Celsius scale or the centigrade scale, is a temperature scale based on 0⁰ for the freezing point of water and 100⁰ for the boiling point of water. This scale was first introduced by Anders Celsius the Swedish physicist.
Why is the Celsius scale more commonly used?
The Celsius scale is more commonly used because it is used along with the metric scale. Since most countries use the metric scale, it makes sense that they also use the Celsius scale. Another reason is that it is easier to convert Celsius into Kelvin, which is another widely used scale for temperatures.
The Fahrenheit scale is a temperature scale based on one proposed in 1724 by the physicist Daniel Gabriel Fahrenheit (1686–1736).
The interval between the ice point and the steam point is divided into 180 equal divisions. Each division is called one degree Fahrenheit and is denoted as °F.
In 1848, Lord Kelvin defined an absolute temperature scale based on the Carnot cycle which was later named after him as Kelvin’s absolute temperature scale. In Kelvin’s scale, the zero point is 273.15 below that of the Celsius scale. The true origin of the universe, if it had one, remains a mystery for the present and likely will remain one far into the future.
Each division in the kelvin scale called a kelvin (K) is equal to a degree on the Celsius scale, but the difference is where zero is. In the Celsius scale 0o is the freezing point of water while in the Kelvin scale the zero point is at absolute zero. Therefore, 0oK is equal to -273.15⁰C, 0⁰C is equal to 273.15 kelvins. The Kelvin scale is used for very low or very high temperatures when water is not involved.
The Kelvin temperature scale is an absolute temperature scale with zero at absolute zero. Because it is an absolute scale, measurements made using the Kelvin scale do not have degrees.
↔ Uncertainty in Measurements
All scientific measurements involve certain degree of error or uncertainty. The errors which arise depend upon two factors.
(i) Skill and accuracy of the worker
(ii) Limitations of measuring instruments. ↔
↔ Scientific Notation
Scientific notation is a form of presenting very large numbers or very small numbers in a simpler form.
It is an exponential notation in which any number can be represented in the form N x 10ⁿ where n is an exponent having positive or negative values and N can vary between 1 to 10.
Important rule for Scientific notation- a. The base is taken as 10. b. The exponent must be a non-zero integer, that means it can be either positive or negative. Exponent is defined as the method of expressing large numbers in terms of powers.
c. The absolute value of the coefficient is greater than or equal to 1 but it should be less than 10. Coefficients can be positive or negative numbers including whole and decimal numbers d. The mantissa carries the rest of the significant digits of the number. What is mantissa and exponent?
In decimal, very large numbers can be shown with a mantissa and an exponent. i.e. 0.12X 10² Here the 0.12 is the mantissa and the 10² is the exponent. the mantissa holds the main digits and the exponents defines where the decimal point should be placed. The same technique can be used for binary numbers.
Thus, 232.508 can be written as 2.32508 x 10² in scientific notation.
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. Now let us see how calculations are carried out with numbers expressed in scientific notation.
(i) Calculation involving multiplication and division ↔
Multiplication and Division
Fig. (031101)
Addition and Subtraction For these two operations, first the numbers are written in such a way that they have the same exponent. After that, the coefficients (digit terms) are added or subtracted as the case may be.
Adding and Subtracting with Scientific Notation - Steps Step 1: Rewrite the numbers so that they all have the same power of ten. Step 2: Add the numbers. Step 3: Rewrite the sentence in a scientific notation.
e.g. Fig. (031101)
↔ Significant Figures
Significant figures are meaningful digits which are known with certainty. There are certain rules for determining the number of significant figures. These are stated below:
a. 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.
b. Zeros preceding to first non-zero digit are not significant. Such zeros indicates the position of decimal point.
e.g. 0.03 has one significant figure and 0.0052 has two significant figures.
c. Zeros between two non-zero digits are significant. Thus, 2.005 has four significant figures.
d. Zeros at the end or right of a number are significant provided they are on the right side of the decimal point. For example, 0.200 g has three significant figures.
e. Counting numbers of objects. 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
Precision refers to the closeness of various measurements for the same quantity. However, accuracy is the agreement of a particular value to the true value of the result.
Precision and accuracy are two ways that scientists think about error.
Accuracy refers to how close a measurement is to the true or accepted value.
Precision refers to how close measurements of the same item are to each other. Precision is independent of accuracy.
e.g.
True value for a result is 2.00 g.
Student ‘A’ takes two measurements and reports the results as 1.95 g and 1.93 g. These values are precise as they are close to each other but are not accurate.
Another student ‘B’ repeats the experiment and obtains 1.94 g and 2.05 g as the results for two measurements. These observations are neither precise nor accurate.
When the third student ‘C’ repeats these measurements and reports 2.01 g and 1.99 g as the result, these values are both precise and accurate.
Table (031101).
Here, 18.0 has only one digit after the decimal point and the result should be reported only up to one digit after the decimal point, which is 31.1. Addition and Subtraction of Significant Figures
In addition or subtraction of the numbers having different precisions, the final result should be reported to the same number of decimal places as in the term having the least number of decimal places.
e.g. Table (031101).
Subtraction of numbers can be done in the same way as the addition.
Multiplication and Division of Significant Figures
In these operations, the result must be reported with no more significant figures as in the measurement with the few significant figures.
2.5×1.25 = 3.125
Since 2.5 has two significant figures, the result should not have more than two significant figures, thus, it is 3.1.
Points for rounding off the numbers 1. 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.
2. 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.
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 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.
Dimensional Analysis
Often while calculating, 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. This is illustrated below.
Example
A piece of metal is 3 inch (represented by in) long. What is its length in cm?
Table (031101).
Dimensional Analysis Often while calculating, 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.
e.g. (031101.01210)

Laws of chemical Combinations

Law of Conservation of Mass
Law of Definite Proportions
Law of Multiple Proportions
Gay Lussac’s Law of Gaseous Volumes
Avogadro’s Law
Dalton’s Atomic Theory
The law of conservation of mass states that the total mass of all the products in a chemical reaction is the same as that of all the reactants present before the reaction.
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Laws of chemical combinations
This law was put forth by Antoine Lavoisier in 1789.
The law of conservation of mass states that
“The mass in an isolated system can neither be created nor be destroyed but can be transformed from one form to another”
According to the law of conservation of mass, the mass of the reactants must be equal to the mass of the products for a low energy thermodynamic process.
Law of Definite Proportions
French chemist Joseph Proust proposed the law of definite composition or proportions based on his experiments.
The law of definite proportions, also known as the law of constant proportions, states ‘that the individual elements that constitute a chemical compound are always present in a fixed ratio’ (in terms of their mass). This ratio does not depend on the source of the chemical compound or the method through which it was prepared."
Proust worked with two samples of cupric carbonate - one of which was of natural origin and the other was synthetic. He found that the composition of elements present in it was same for both the samples as shown below:

Thus, he concluded that irrespective of the source, a given compound always contains same elements combined together in the same proportion by mass. This law sometimes also referred to as Law of Definite Composition

The law of multiple proportions states that when two elements combine to form two or more compounds, then masses of one of the elements which combine with a fixed mass of the other bear a simple ratio to one another.
The Gay-Lussac’s law of combining volumes states that, when different gases react together, they do so in terms of volume which bears simple whole-number ratio provided that the temperature and pressure of reacting gases remain constant.
e.g.
H₂ (g) + Cl₂(g)→ 2HCl
If we take two volumes of hydrogen and one volume of chlorine, both would react with each other to form two volumes of hydrogen chloride.
Based on Gay-Lussac's results, Avogadro gave his hypotheses that, at the same pressure and temperature, equal volumes of gas contain equal numbers of molecules (Avogadro’s Law). 2 molecules of hydrogen + 1 molecule of chlorine= 2 molecules of hydrogen chloride It can also be expressed in a different way, for example, consider 100 mL of hydrogen combined with 100 mL of chlorine to give 200 mL of hydrogen chloride.
Hydrogen (100 mL) + Chlorine (100 mL) = Hydrogen Chloride (200 mL)
Thus, the volumes of hydrogen and chlorine which combine (i.e., 100mL and 100mL) bear a simple ratio of 1:1 and thus the simple ratio volumes are 1:1:2.

Avogadro’s law states that equal volumes of all gases under similar conditions of temperature and pressure contain equal numbers of molecules.

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

Atomic and Molecular Masses Atomic Mass
The atomic mass of an element is the number of times an atom of that element is heavier than an atom of carbon taken as 12.
It may be noted that the atomic masses as obtained above are the relative atomic masses and not the actual masses of the atoms.
One atomic mass unit (amu) is equal to l/12th of the mass of an atom of carbon-12 isotope. It is also known as unified mass.

Average Atomic Mass
Most of the elements exist as isotopes which are different atoms of the same element with different mass numbers and the same atomic number.
Therefore, the atomic mass of an element must be its average atomic mass and it may be defined as the average relative mass of an atom of an element as compared to the mass of carbon atoms (C-12) taken as 12u.

Molecular Mass
Molecular mass is the sum of atomic masses of the elements present in a molecule. It is obtained by multiplying the atomic mass of each element by number of its atoms and adding them together.
e.g. Molecular mass of methane (CH₄)
= 12.011 u + 4 (1.008 u)
= 16.043 u

Formula Mass
Ionic compounds such as NaCl, KNO₃, Na₂CO₃ etc. do not consist of molecules i.e., single entities but exist “as ions closely packed together in a three dimensional space as shown in –Fig. In such cases, the formula is used to calculate the formula mass instead of molecular mass.
Thus, formula mass of NaCl = Atomic mass of sodium + atomic mass of chlorine
= 23.0 u + 35.5 u = 58.5 u. ↔

Mole Concept
The mole, symbol mol, is the SI unit of amount of substance. One mole contains exactly 6.02214076 × 10²³ elementary entities. This number is the fixed numerical value of the Avogadro constant, NA, when expressed in the unit mol–1 and is called the Avogadro number.
The amount of substance, symbol n, of a system is a measure of the number of specified elementary entities. An elementary entity may be an atom, a molecule, an ion, an electron, any other particle or specified group of particles.
It may be emphasised that the mole of a substance always contains the same number of entities, no matter what the substance may be.
In order to determine this number precisely, the mass of a carbon–12 atom was determined by a mass spectrometer and found to be equal to 1.992648 × 10^-23 g. Knowing that one mole of carbon weighs 12 g, the number of atoms in it is equal to:

=6.0221367X10²³ atoms/mole.
We can, therefore, say that 1 mol of hydrogen atoms = 6.022 × 1023 atoms
1 mol of water molecules = 6.022 × ²³ water Molecules
1 mol of sodium chloride = 6.022 ײ³ formula units of sodium chloride.
The mass of one mole of a substance in grams is called its molar mass. The molar mass in grams is numerically equal to atomic/molecular/ formula mass in u. Molar mass of water = 18.02 g mol^-1 Molar mass of sodium chloride = 58.5 g mol^-1

Percentage Composition One can check the purity of a given sample by analysing this data. Let us understand by taking the example of water (H₂O). Since water contains hydrogen and oxygen, the percentage composition of both these elements can be calculated.
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Empirical Formula The formula of the compound which gives the simplest whole number ratio of the atoms of yarious elements present in one molecule of the compound. For example, the formula of hydrogen peroxide is H₂O₂. In order to express its empirical formula, we have to take out a common factor 2. The simplest whole number ratio of the atoms is 1:1 and the empirical formula is HO. Similarly, the formula of glucose is C₆H₁₂O₆. In order to get the simplest whole number of the atoms, Common factor = 6 The ratio is = 1 : 2 : 1 The empirical formula of glucose = CH₂O
Molecular Formula
The formula of a compound which gives the actual ratio of the atoms of various elements present in one molecule of the compound. For example, molecular formula of hydrogen peroxide = H₂O₂ and Glucose = C₆H₁₂O₆.
Molecular formula = n x Empirical formula
Where n is the common factor and also called multiplying factor. The value of n may be 1, 2, 3, 4, 5, 6 etc.
In case n is 1, Molecular formula of a compound = Empirical formula of the compound.
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Stoichiometry and Stoichiometric Calculations 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 volume also) of the reactants and the products involved in a chemical reaction. ↔
Let us consider the combustion of methane. A balanced equation for this reaction is as given below:
→ Limiting Reactant/Reagent Sometimes, in alchemical equation, the reactants present are not the amount as required according to the balanced equation. The amount of products formed then depends upon the reactant which has reacted completely. This reactant which reacts completely in the reaction is called the limiting reactant or limiting reagent. The reactant which is not consumed completely in the reaction is called excess reactant.
Reactions in Solutions
When the reactions are carried out in solutions, the amount of substance present in its given volume can be expressed in any of the following ways:
1. Mass percent or weight percent (w/w%)
2. Mole fraction
3. Molarity
4. Molality
1. Mass percent: It is obtained by using the following relation: 2. Mole fraction: It is the ratio of number of moles of a particular component to the total number of moles of the solution.
3. Molarity: It is defined as the number of moles of solute in 1 litre of the solution.
4. Molality: It is defined as the number of moles of solute present in 1 kg of solvent. It is denoted by m. ↔
1. AZT is used for the victims of (AIDS).
2. The prefix pico stands for (10-12).
3. Pascal are the units for the physical quantity (Pressure).
4. The number of significant frgures in 0.00030 is (two)
6. Decimal equivalent of 2/3 is upto three significant figures (.667).
6. The empirical formula of hydrogen peroxide is (HO).
7. The law which does not follow from Dalton's atomic theory is (Gay-Lussac law of gaseous volumes)
8. The mass of a molecule of carbon-l4 dioxide (14CO2) is (7.64X10-23).
9. An atom of sulphur is (8/3) times heavier than an atom of carbon.
10. (1.9) mol of N, are needed to produce 3.8 mol of NH, by reaction with hydrogen.
11. If mole fraction ofsodium chloride in sodium chloride aqueous solution is 0.35, then mole fraction of water in the solution is (.65).
12. Amount of glucose (C6H12O6) required to prepare 100 mL of 0.1 M solution is (1.8g).
13. The empirical formula of benzene is (CH).
SVTs
SVT031101.001

1. Name a device which is used to measure force acting on an object.

Solution Dynamometer is the device used for measuring force acting on an object
2. What is the formula to convert from Celsius to Fahrenheit?
The formula to convert from Celsius to Fahrenheit is F= (9/5 x C) + 32.
3. What is the formula to convert from Fahrenheit to Celsius?
The formula to convert from Fahrenheit to Celsius is C = 5/9(F – 32).
4. What is the SI unit of temperature?
4. The SI unit of temperature is Kelvin.
5. What are the other units of temperature?
Other units of temperature are Rømer, Rankine, Réaumur, Delisle, and Newton. 5. What are the three common temperature scales?
The three common temperature scales are
Celsius, Fahrenheit and Kelvin
6. What is Density?
A material’s density is defined as its mass per unit volume.
7. Who discovered the principle of Density?
The principle of density was discovered by the Greek scientist Archimedes.
8. How would you find the density of a human body?
The density of a human body can be determined by the following expression:
Density = Mass / Volume
The mass of the human body can be calculated by using a weight scale. The submersion displacement is used to get the volume of the human body. If we fill a tub with water and let the person fully submerge into the water, then we see a rise in the water, this rise is equal to the volume of the body. Thus, the density of the human body can be obtained by dividing the mass of the human body by the volume of the human body.
9. How will we know if a substance is less dense than water?
If a substance weighs less than an equal volume of water, it is less dense and will float.
10. What happens to the least dense of two immiscible liquids?
Immiscible liquids: Liquids that don't mix or are not soluble in each other, are called immiscible liquids. For example petrol and water are two immiscible liquids. Petrol and water are an example of immiscible liquids so they are not soluble to each other, e.g. Kerosene and water, oil and water, benzene and water, honey and oil, etc.
Miscible liquids are liquids that dissolve in each other to form a homogeneous mixture while immisible liquids are those liquids that do not mix well in each other. Alcohol and water are miscible liquids while oil and water are immiscible pair of liquids, e.g. Acetic Acid and Water, Gasoline (Petrol) and Deisel, Milk Coffee. Lemonade etc.
If the liquids are immiscible, and they are not stirred, or only stirred gently, they will separate into 2 layers, with the less dense floating on the more dense liquid. There will be a small amount of dissolving at the interface, but this will quickly reach equilibrium, with very small concentrations of each liquid dissolved in the other. The process is called decantation.

E1. A piece of metal is 3 inch (represented by in) long. What is its length in cm?
E2. A jug contains 2 L of milk. Calculate the volume of the milk in m³.
E3. How many seconds are there in 2 days?
E4.1. Calculate the molecular mass of glucose (C6H12O6) molecule.
B1. Calculate the molar mass of the following:
(i) H₂O (ii) CO₂ (iii) CH₄
E5.2. A compound contains 4.07% hydrogen, 24.27% carbon and 71.65% chlorine. Its molar mass is 98.96 g. What are its empirical and molecular formulas?
E6.3. Calculate the amount of water (g) produced by the combustion of 16 g of methane.
E7.4. How many moles of methane are required to produce 22g CO₂ (g) after combustion?

E8.5. 50.0 kg of N2 (g) and 10.0 kg of H2 (g) are mixed to produce NH3 (g). Calculate the amount of NH₃ (g) formed. Identify the limiting reagent in the production of NH₃ in this situation.
E9.6. A solution is prepared by adding 2 g of a substance A to 18 g of water. Calculate the mass per cent of the solute.
E10.7. Calculate the molarity of NaOH in the solution prepared by dissolving its 4 g in enough water to form 250 mL of the solution.
E11.8. The density of 3 M solution of NaCl is 1.25 g mL–1. Calculate the molality of the solution.

031101.01210
B2. Calculate the mass per cent of different elements present in sodium sulphate (Na₂SO₄).
B3. Determine the empirical formula of an oxide of iron, which has 69.9% iron and 30.1% dioxygen by mass.
B4. Calculate the amount of carbon dioxide that could be produced when (i) 1 mole of carbon is burnt in air.
(ii) 1 mole of carbon is burnt in 16 g of dioxygen.
(iii) 2 moles of carbon are burnt in 16 g of dioxygen.
B5. Calculate the mass of sodium acetate (CH₃COONa) required to make 500 mL of 0.375 molar aqueous solution. Molar mass of sodium acetate is 82.0245 g mol-1.
B6. Calculate the concentration of nitric acid in moles per litre in a sample which has a density, 1.41 g mL^-1and the mass per cent of nitric acid in it being 69%.
B7. How much copper can be obtained from 100 g of copper sulphate (CuSO₄)?
B8. Determine the molecular formula of an oxide of iron, in which the mass per cent of iron and oxygen are 69.9 and 30.1, respectively.
B9. Calculate the atomic mass (average) of chlorine using the following data:

B10. In three moles of ethane (C₂H₆), calculate the following:
(i) Number of moles of carbon atoms.
(ii) Number of moles of hydrogen atoms.
(iii) Number of molecules of ethane.

031101.11220
B11. What is the concentration of sugar (C₁₂H₂₂O₁₁) in mol L^-1if its 20 g are dissolved in enough water to make a final volume up to 2L?
B12. If the density of methanol is 0.793 kg L^-1, what is its volume needed for making 2.5 L of its 0.25 M solution?
B13. Pressure is determined as force per unit area of the surface. The SI unit of pressure, pascal is as shown below:
1Pa = 1N m^-2
B14. What is the SI unit of mass? How is it defined?

B15. Match the following prefixes with their multiples:

B16. What do you mean by significant figures?
B17. A sample of drinking water was found to be severely contaminated with chloroform, CHCl³, supposed to be carcinogenic in nature. The level of contamination was 15 ppm (by mass).
(i) Express this in per cent by mass.
(ii) Determine the molality of chloroform in the water sample.
B18. Express the following in the scientific notation:
(i) 0.0048
(ii) 234,000
(iii) 8008
(iv) 500.0
(v) 6.0012
B19. How many significant figures are present in the following?
(i) 0.0025
(ii) 208
(iii) 5005
(iv) 126,000
(v) 500.0
(vi) 2.0034
B20. Round up the following upto three significant figures:
(i) 34.216
(ii) 10.4107
(iii) 0.04597
(iv) 2808
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B21. The following data are obtained when dinitrogen and dioxygen react together to form different compounds:

(a) Which law of chemical combination is obeyed by the above experimental data? Give its statement.
(b) Fill in the blanks in the following conversions:
(i) 1 km = ...................... mm = ...................... pm
(ii) 1 mg = ...................... kg = ...................... ng
(iii) 1 mL = ...................... L = ...................... dm³
B22. If the speed of light is 3.0 × 10⁸ m s^-1, calculate the distance covered by light in 2.00 ns.
B23. In a reaction
A + B₂→AB₂

Identify the limiting reagent, if any, in the following reaction mixtures.
(i) 300 atoms of A + 200 molecules of B
(ii) 2 mol A + 3 mol B
(iii) 100 atoms of A + 100 molecules of B
(iv) 5 mol A + 2.5 mol B
(v) 2.5 mol A + 5 mol B
B24. Dinitrogen and dihydrogen react with each other to produce ammonia according to the following chemical equation:
N₂ (g) + H₂ (g) → 2NH₃ (g) (i) Calculate the mass of ammonia produced if 2.00 × 10³ g dinitrogen reacts with 1.00 × 10³ g of dihydrogen.
(ii) Will any of the two reactants remain unreacted?
(iii) If yes, which one and what would be its mass?
B25. How are 0.50 mol Na₂CO₃ and 0.50 M Na₂CO₃ different?
B26. If 10 volumes of dihydrogen gas reacts with five volumes of dioxygen gas, how many volumes of water vapour would be produced?
B27. Convert the following into basic units:
(i) 28.7 pm
(ii) 15.15 pm
(iii) 25365 mg
B28. Which one of the following will have the largest number of atoms?
(i) 1 g Au (s)
(ii) 1 g Na (s)
(iii) 1 g Li (s)
(iv) 1 g of Cl(g)
B29. Calculate the molarity of a solution of ethanol in water, in which the mole fraction of ethanol is 0.040 (assume the density of water to be one).
B30. What will be the mass of one ¹²C atom in g?
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B31. How many significant figures should be present in the answer of the following calculations?
(i)

(ii) 5 × 5.364
(iii) 0.0125 + 0.7864 + 0.0215
B32. Use the data given in the following table to calculate the molar mass of naturally occurring argon isotopes:

B33. Calculate the number of atoms in each of the following (i) 52 moles of Ar (ii) 52 u of He (iii) 52 g of He.
B34. A welding fuel gas contains carbon and hydrogen only. Burning a small sample of it in oxygen gives 3.38 g carbon dioxide, 0.690 g of water and no other products. A volume of 10.0 L (measured at STP) of this welding gas is found to weigh 11.6 g. Calculate (i) empirical formula, (ii) molar mass of the gas, and (iii) molecular formula.
B35. Calcium carbonate reacts with aqueous HCl to give CaCl₂ and CO₂ according to the reaction,
CaCO₃ (s) + 2 HCl (aq) → CaCl₂ (aq) + CO₂ (g) + H₂O(l)
What mass of CaCO₃ is required to react completely with 25 mL of 0.75 M HCl?
B36 Chlorine is prepared in the laboratory by treating manganese dioxide (MnO2) with aqueous hydrochloric acid according to the reaction 4 HCl (aq) + MnO₂ (s) → 2H₂O (l) + MnCl₂ (aq) + Cl₂ (g) How many grams of HCl react with 5.0 g of manganese dioxide?
MCQs
1. A pure substance which contains only one type of atom is called ————–.
(a) An element (b) a compound (c) a solid (d) a liquid

(a) An element
Solution: An element is made up of only one type of atom.

2. Two different oxides of a metal contain 20% and 27% oxygen by weight. This is in accordance with the law of:
(a) conservation of mass (b) constant composition (c) multiple proportion (d) reciprocal proportion.

C

3. The molar masses of oxygen and sulphur dioxide are 32 and 64 respectively. if 1 L of oxygen at 25⁰C and. 750 mm Hg pressure contains N molecules, then the number of rnolecules in 2 L sulphur dioxide under same conditions of temperature and pressure is :
(a) N/2 (b) 3N/2 (c) 2N (d) 6 N.

C According to avogardo hypothesis , equal volumes of all gases under similar conditions of temperature and pressure contain equal number of molecules . Therefore ,if 1L of O2(or SO2 ) contains N molecules at 25∘
C and 750 mm Hg pressure , then 2L of O2 (or SO2 ) contain 2N molecules at same conditions of temperature and pressure.

4. 2 g of oxygen contain number of atoms equal to that contained by : (a) 0.5 g hydrogen (b) 4 g sulphur (c) 7 g nitrogen (d) 2.3 g sodium.

B

5. The smallest particle that can take part in chemical reactions is ————–.
(a) Atom (b) molecule (c) Both (a) and (b) (d) none of these

A Solution: The smallest particle that can take part in chemical reactions is both an atom.

6. Which of the following is a homogeneous mixture? (a) Mixture of soil and water (b) Sugar solution (c) Mixture of sugar, salt and sand (d) Iodised table salt

(b) Sugar solution Solution: Sugar solution is a homogeneous mixture. A homogeneous mixture is a mixture in which the composition is uniform throughout the mixture

7. The significant figures in 0.00051 are ————–.
(a) 5 (b) 3 (c) 2 (d) 26

(c) 2 Solution: The significant figures in 0.00051 are 2.

8. Formation of CO and Carbon dioxide illustrates the law of ————–.
(a) Law of conservation of mass (b) Law of Reciprocal proportion (c) Law of Constant Proportion (d) Law of Multiple Proportion

(d) Law of Multiple Proportion Solution: If an element forms more than one compound with another element for a given mass of an element, masses of other elements are in the ratio of small whole numbers.

9. The number of significant figures in 6.02 x 1023 is ————–.
(a) 23 (b) 3 (c) 4 (d) 26

(b) 3 Solution: The number of significant figures in 6.02 x 1023 is 3

10. The prefix 1018 is ————–. (a) giga (b) exa (c) kilo (d) mega

(b) exa Solution: The prefix 1018 is exa

11. The mass of an atom of carbon is ————–.
(a) 1g (b) 1.99 x 10-23 g (c) 1/12 g (d) 1.99 x 1023 g

(b) 1.99 x 10 -23 g Solution: The mass of an atom of carbon is {12 / (6.02 x 10 23)} = 1.99 x 10 -23 g

12. A measured temperature on Fahrenheit scale is 200F. What will this reading be on the Celsius Scale?
(a) 40 ℃ (b) 94 ℃ (c) 93.3 ℃ (d) 30 ℃

(c) 93.3 ℃ Solution: The relationship between Fahrenheit and degree Celsius is: (0F) = 9/5 (oC) +32.

13. Which of the following pairs of gases contains the same number of molecules? (a) 16 g of O2 and 14 g of N2 (b) 6 g of O2 and 22 g of CO2 (c) 28 g of N2 and 22 g of CO2 (d) 32 g of CO2 and 32g of N2

(a) 16 g of O2 and 14 g of N2 Solution: Divide the given mass by its molar mass to get moles, then multiply times 6.022×1023 to get the number of molecules.

14. Device used to measure the force acting as an object is a. spring balance b. measuring tape c. weighing scale d. electrometer

A Spring balance is a device used to measure the force acting on an object. The stretch in the spring gives the magnitude of the force.

15. Device used to measure the force acting as an object is
a. spring balance b. measuring tape c. weighing scale d. electrometer

A spring balance Spring balance is a device used to measure the force acting on an object. The stretch in the spring gives the magnitude of the force.

16. The device used to measure the relative density of a liquid is known as: a. Thermometer b. Hydrometer c. Barometer d. Spring balance

B Hydrometer A hydrometer is an instrument used to measure the relative density of liquids. It is usually made of glass and consists of a cylindrical stem and a bulb weighted with mercury or lead shot to make it float upright.

17.
1. Separate the following cations into monovalent, divalent and trivalent cations. Also write their formulae with respective valencies. hydrogen ion; lithium ion; sodium ion; potassium ion; beryllium ion; magnesium ion; calcium ion; barium ion; aluminum ion; silver ion; zinc ion; cuprous ion; cupric ion; ferrous ion; ferric ion; stannous ion; stannic ion; auric ion; mercurous ion; mercuric ion; aurous ion; plumbous ion and plumbic ion.
2. Separate the following anions into monatomic and polyatomic ions (divalent cation and trivalent cation. Also write their formulae with respective valencies.
Hydride; phosphate; thiosulfate; hydrogen sulfate (bisulfate); hydrogen sulfite (bisulfite); fluoride; chloride; bromide; iodide; oxide; sulfide; nitride; phosphide; nitrate; nitrite; chromate; dichromate; cyanide; permanganate; hydroxide; carbonate; hydrogen carbonate (bicarbonate); sulfate; sulfite; oxalate and phosphate Hint- A monovalent cation has unit positive charge whereas a monovalent anion has unit negative charge. A divalent cation has two positive charges whereas a divalent anion has two negative charges. A trivalent cation has three positive charges whereas a trivalent anion has three negative charges.
Symbol Name Symbol Name H+ hydrogen ion H- hydride Li+ lithium ion F- fluoride Na+ sodium ion Cl- chloride K+ potassium ion Br- bromide Rb+ rubidium ion I- iodide Cs+ cesium ion O2- oxide Be2+ beryllium ion S2- sulfide Mg2+ magnesium ion Se2- selenide Ca2+ calcium ion Te2- telluride Sr2+ strontium ion N3- nitride Ba2+ barium ion P3- phosphide Ra2+ radium ion As3- arsenide Ag+ silver ion Zn2+ zinc ion Al3+ aluminum ion Systematic name Common Systematic name Common Symbol (Stock system) name Symbol (Stock system) name Cu+ copper(I) cuprous Hg22+ mercury(I) mercurous Cu2+ copper(II) cupric Hg2+ mercury(II) mercuric Fe2+ iron(II) ferrous Pb2+ lead(II) plumbous Fe3+ iron(III) ferric Pb4+ lead(IV) plumbic Sn2+ tin(II) stannous Co2+ cobalt(II) cobaltous Sn4+ tin(IV) stannic Co3+ cobalt(III) cobaltic Cr2+ chromium(II) chromous Ni2+ nickel(II) nickelous Cr3+ chromium(III) chromic Ni4+ nickel(IV) nickelic Mn2+ manganese(II) manganous Au+ gold(I) aurous Mn3+ manganese(III) manganic Au3+ gold(III) auric Symbols and Charges for Polyatomic Ions Formula Name Formula Name NO3- nitrate ClO4- perchlorate NO2- nitrite ClO3- chlorate CrO42- chromate ClO2- chlorite Cr2O72- dichromate ClO- hypochlorite CN- cyanide IO4- periodate MnO4- permanganate IO3- iodate OH- hydroxide IO- hypoiodite O22- peroxide BrO3- bromate NH2- amide BrO- hypobromite CO32- carbonate HCO3- hydrogen carbonate (bicarbonate) SO42- sulfate HSO4- hydrogen sulfate (bisulfate) SO32- sulfite HSO3- hydrogen sulfite (bisulfite) C2O42- oxalate HC2O4- hydrogen oxalate (binoxalate) PO43- phosphate HPO42- hydrogen phosphate PO33- phosphite H2PO4- dihydrogen phosphate S2O32- thiosulfate HS- hydrogen sulfide AsO43- arsenate BO33- borate SeO42- selenate B4O72- tetraborate SiO32- silicate SiF62- hexafluorosilicate C4H4O62- tartrate