Many phenomenon like formation of milk form curd, formation of vinegar form for sugarcane juice(ganne ka juice), rusting of iron etc. are some examples of chemical changes in our daily life. The branch of science that studies the composition, structure, properties, and changes of matter is called chemistry.
Chemistry mainly develop in form of alchemy and latrochemistry during 1300-1600 CE. The 18th century in Europe, the modern chemistry took shape and the alchemical tradition introduced in Europe by Arabs. In ancient India, the chemistry was also called Rasayan Shastra, Ras kriya, Ras vidya and Rastantra.
Chemistry is most important in out daily life such as many drugs like cisplatin and taxol which used in treatment of cancer and ATZ(azidothymidine) are used in AIDS treatment, etc.
MATTER-
- Matter- Anything which have any mass and occupy space(volume).
- Different types(states) of matter-
| State of mater | Arrangement of particles | Shape | Volume | Compressibility | Examples |
| Solid | Particles are tightly packed in a fixed arrangement with strong intermolecular forces. | Definite | Definite | Negligible; particles cannot be compressed easily. | Wood, ice, iron, etc. |
| Liquid | Particles are close but not in a fixed position; they can flow and slide past each other. | Indefinite; takes the shape of the container where it placed | Definite | Slight; particles can move slightly closer under pressure | Oil, water, milk, etc. |
| Gas | Particles are far apart with negligible intermolecular forces and move freely in all directions. | Indefinite; spreads to fill the container. | Indefinite; expands to fill the container. | Highly compressible due to large spaces between particles. | Oxygen, Carbon dioxide, Steam, etc. |
- Further classification of matter- Matter are further classify as mixture and pure substance.
- Mixture-
- It contain particles of 2 or more pure substance which may be present in any ratio in it. ex-air, tea, sugar in water etc.
- Mixture have 2 types- Homogenous and heterogenous
- Homogenous Mixture- It is a uniform combination of 2 or more substances where components are indistinguishable and evenly distributed. ex- air, sugar solution etc.
- Heterogenous Mixture- It is a non-uniform combination of 2 or more substance where components are distinguishable and unevenly distributed and some time we can see the particle of different components. ex- sand in water etc.
- Pure Substance-
- It contain elements of single atom and have fixed composition. ex-copper, gold etc.
- Pure further classify into 2- elements and compound.
- Element- It contain only single type of atoms. It particles may in form of atoms or molecules. ex- Cu, Al, Zn etc.
- Compound- It contain atoms of different elements in definite ratio or fixed composition. Compound cannot be separated by physical method.
- Interesting note- When hydrogen burn, then pop sound released where oxygen support this combustion but water work as fire extinguisher.
- Mixture-
- Physical properties of matter-
- Those properties which can measure without changing the identity of substance. ex. color, odor, melting point, boiling point, density, state, conductivity, volume etc.
- Chemical properties of matter-
- Those properties which measure after changing the identity or composition of substance. ex. combustion, acidity, Corrosion, etc.
- System of measurement matter-
- Earlier, mainly 2 different system used in measurement- English System and Metric System (originate in France in late 18th century) in different regions in the world.
- But the scientist community face many problems regarding measurement in different areas of world, so they need to standarise the system of measurement then SI system introduce.
- SI Unit-
- The full form of SI in SI unit is “Système International d’Unités”, which is French for “International System of Units.
- It stablished by 11th General conference of weight and measure.
- The CGPM is intergovernmental treaty organization which created by diplomatic treaty known as Metre Convention signed in Paris in 1857.
- The SI system has 7 base unit-
| Base Physical Qyantity | Symbol of Quantity | Name of SI Unit | Symbol of SI Unit |
|---|---|---|---|
| Length | l | Meter | m |
| Mass | m | Kilogram | kg |
| Time | t | Second | s |
| Electric Current | I | Ampere | A |
| Thermodynamic Temperature | T | Kelvin | K |
| Amount of Substance | n | Mole | mol |
| Luminous Intensity | lv | Candela | cd |
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Some Important Physical Quantities-
- Mass-
- It is the amount of matter in an object.
- Its SI unit is Kg.
- It is scalar quantity.
- It is constant quantity because it doesn’t change with location.
- It is accurately determine by analytical balance, scale in laboratory.
- Weight-
- It is the force exerted by gravity on an object.
- Its SI unit is Newton(N).by
- It is varies with place to place, weight changes with the gravitational field.
- Weight(W) = Mass(M) * Gravity(G)
- It is vector quantity.
- It is also measure from Spring balance.
- Volume-
- It is a amount of space which occupy by object.
- Its unit is (length)3, so SI unit is (meter)3.
- A common unit used is liter(L) which is not SI unit but mostly used.
- 1 L = 1000ml, 1 dm3 = 1000 cm3.
- It is measure through Volumetric flask, Graduated cylinder, Pipette, Burette.
- Density(ρ)-
- It is a amount of mass as per unit volume.
- Density = Mass / Volume
- Its SI unit is Kg/m3 or Kg m-3.
- Its tells that how closely the particles are packed.
- It is scalar quantity.
- Temperature-
- It measure that how objects are hot and cold.
- It SI unit is Kelvin(K).
- Also Celsius(°C) and Fahrenheit(°F) are used to calculate temperature.
- Some important formulas-
- °F = 9/5(°C) + 32
- °C = 5/9(°F) – 32
- K = °C + 273.15
- It is scalar quantity.
- It is commonly measure through thermometer.
- Some important values of temperature-
- Human body temperature- 37°C or 98.6°F or 310 K.
- Room temperature- 25°C or 77°F or 298 K.
- Boiling point- 100°C or 212°F or 373.15 K.
- Freezing point- 0°C or 32°F or 273.15 K.
Uncertainty In Measurement-
- Significant figures and its rule-
- It is the digits in a measurement that are important for precision.
- Rules to determine significant figures-
- Non-zero digits are always significant. ex.- 345 has 3 significant figures.
- Zeros between non-zero digits are significant. ex.- 1002 has 4 significant figures.
- Leading zeros (zeros before non-zero digits) are not significant.
- Trailing zeros in a decimal number are significant. ex.- 0.00456 has 3 significant figures. ex.- 45.00 has 4 significant figures.
- Trailing zeros in a whole number without a decimal point are not significant. ex.- 500 has 1 significant figure.
- Exact numbers (like counted items or constants) have infinite significant figures. ex.- 1 dozen = 12 (infinite significant figures).
- Rules of addition/subtraction and multiplication/division are-
- Addition/Subtraction: The result should have the same number of decimal places as the number with the least decimal places.
- Multiplication/Division: The result should have the same number of significant figures as the number with the least significant figures.
- Scientific notation-
- It is a way of expressing very large or very small numbers in the form of a * 10n. (ex.- 5000 can be written as 5 * 103, 0.00045 can be written as 4.5 * 10-4, etc.).
- Dimension analysis-
- It is a way of conversion of one units to another unit. [You read more in physics chapter 1].
Law of Chemical Combination-
- There are 5 basic law to combine elements to form compounds-
- Law of definite proportion(by French chemist, Joseph Proust in 1799)- It states that a compound always contains the same elements in the same fixed ratio by mass, regardless of its source or amount.
- Law of multiple proportion(by Dalton in 1803)– It state that the two elements combine to form different compounds, then the ratios of their masses are in simple whole numbers.
- Law of Conservation of mass(by Antonine Lavoisier in 1789)– It state that the mass can neither be created nor be destroy, it only transfer form one form to another.
- Gay Lussac ‘s law of gaseous volume(by Gay Lussac in 1808)- It states that gases react in simple whole number ratios of their volumes, provided the conditions of temperature and pressure remain constant.
- Avogadro ‘s law(by Avogadro in 1811)- It states that equal volumes of gases at the same temperature and pressure contain equal numbers of molecules.
Some Important Concepts-
- Dalton atomic theory(in )-
- It state that-
- Matter is made of tiny particles called atoms.
- Atoms of the same element are different atom and different elements have different atoms.
- Atoms cannot be created, destroyed, or divided.
- Atoms combine in fixed whole-number ratios to form compounds.
- In chemical reactions, atoms rearrange but remain unchanged.
- Its limitation is-
- It could not explain the law of chemical combination.
- It could not explain the laws of gaseous volumes.
- Also it could not provide the reason of combination of atom.
- Atomic mass-
- Atomic mass is the average mass of an atom of an element, measured relative to 1/12th the mass of a carbon-12 atom. It includes protons and neutrons in the nucleus and is expressed in atomic mass units (amu).
- Molecular mass-
- It is the sum of the atomic masses of all the atoms in a molecule. It is measured in atomic mass units (amu) and represents the mass of a single molecule.
- Average atomic mass-
- It is the weighted average of the masses of all naturally occurring isotopes of an element, based on their abundance. It is measured in atomic mass units (amu) and represents the typical mass of an atom of the element.
- Molecular mass-
- It is the total mass of all the atoms in a molecule, measured in atomic mass units (amu). It is calculated by adding up the atomic masses of each atom in the molecule.
- Formula mass-
- It is the sum of the atomic masses of all atoms in a chemical formula, whether the substance is a molecule or an ionic compound. It is measured in atomic mass units (amu).
- Mole concept-
- The mole concept is a way to count particles (atoms, molecules, or ions) in a substance. One mole equals 6.022 * 1023 particles, called Avogadro’s number.
- It helps relate the mass of a substance to the number of particles it contains.
- The molar mass of a substance in grams is equal to its formula mass in atomic mass units (amu), and one mole of a substance weighs its molar mass in grams.
- No. of moles = given mass or volume / molar mass or volume
- Molar mass-
- It is the mass of one mole of a substance, measured in grams per mole (g/mol). It is equal to the sum of the atomic masses of all atoms in a molecule or formula unit.
Percentage Composition of Compounds-
- Empirical formula-
- It is the simplest whole-number ratio of elements in a compound. It shows the relative number of atoms of each element, not the actual number. For example, the empirical formula of glucose (C₆H₁₂O₆) is CH₂O.
- Rules to find empirical formula is-
- 1. Convert mass percent to grams.
- 2. Convert the mass into moles of each element(moles = given mass / molar mass).
- 3. Find the simplest ratio of moles by dividing each element’s moles by the smallest number of moles.
- 4. Round the ratio to the nearest whole number if necessary.
- 5. Write the empirical formula using the whole-number ratios by multiple n.
- Molecular formula-
- It is original formula which tell actual number of atoms.
- Rules to find molecular formula-
- 1. Determine the empirical formula.
- 2.Find the molar mass of the empirical formula.
- 3. Divide the molecular mass (known or measured) by the molar mass of the empirical formula.
- 4. Multiply the empirical formula by the result to get the molecular formula.
- Rules to find molecular formula-
- n = molecular formula mass / empirical formula mass
- Stoichiometry and its calculations- Stoichiometry is the study of the relationship between the amounts of reactants and products in a chemical reaction. It uses balanced chemical equations to calculate the quantities of substances involved.
- Limiting reagent- It ends or consume first. The limiting reagent is the substance in a chemical reaction that is completely used up first, limiting the amount of product that can be formed.
- It determines how much product can be made.
- Once the limiting reagent is consumed, the reaction stops, even if other reactants are still available.
- Which element remains, called axis reagent.
- Reaction in solution- The concentration of solution is mainly expressed by 4 ways-
- Mass percentage-
- It is the ratio of the mass of a component to the total mass of the solution, multiplied by 100.
- It shows how much of a substance is present in a mixture.
- Mass percent = mass of solute / mass of solutions * 100
- Mole fraction-
- Mole fraction is the ratio of the number of moles of a component to the total number of moles in a mixture.
- It shows the proportion of each substance in a solution.
- Mole fraction = No. of moles(suppose substance a) / No. of mole of solution
- Molarity-
- It measure the concentration of a solution.
- It is the number of moles of solute dissolved in one liter of solution.
- It tells you how strong or concentrated a solution is.
- It is denoted by M.
- M = No. of moles of solute / Volume of solution in liters
- Molality-
- It measure the concentration based on the mass of the solvent.
- It is the number of moles of solute dissolved in one kilogram of solvent.
- It is used when temperature changes because it doesn’t depend on the volume of the solution.
- It is denoted by m.
- m = Moles of solute / Mass of solvent in kilograms
- Normality-
- It measure the concentration of a solution based on its reactive capacity.
- It is the number of gram equivalents of solute per liter of solution
- It is denoted by N.
- N = Gram equivalents of solute / Volume of solution in liters
- Mass percentage-
These all are the notes of chapter 1. After some time you get important question, NCERT solutions HERE. *#THANKS FOR VISITING, VISIT AGAIN#* 😊