Intro to Chemistry Basics

Branches of Chemistry

• Organic Chemistry: deals with the study of carbon compounds, especially hydrocarbons and their derivatives
• Inorganic Chemistry: deals with the study of compounds of all elements except carbon, largely concerns minerals found in the Earth's crust
• Physical Chemistry: explains fundamental principles governing various chemical phenomena
• Industrial Chemistry: studies industrial processes
• Analytical Chemistry: deals with qualitative and quantitative analysis of various substances
• Biochemistry: deals with chemical changes in living organisms
• Nuclear Chemistry: studies nuclear reactions, nuclear fission, nuclear fusion, and transmutation processes

Properties of Matter and Their Measurement

• Properties of matter: classified into physical and chemical properties
• Physical properties: can be measured or observed without changing the identity or composition of the substance (e.g., color, odor, melting point, boiling point, density)
• Chemical properties: characteristic reactions of different substances (e.g., acidity or basicity, combustibility)
• Metric System: based on the decimal system
• International System of Units (SI): established by the 11th General Conference on Weights and Measures (CGPM)
• SI system has seven base units

Prefixes in SI system

• Multiple and prefix symbols (e.g., pico, nano, micro, milli, centi, deci, deca, hecto, kilo, mega, giga, tera)

Mass and Weight

• Mass: amount of matter present in a substance, constant
• Weight: force exerted by gravity on an object, can vary due to change in gravity
• Mass can be determined very accurately using an analytical balance

Volume and Density

• Volume: units of (length)^3, can be measured in m^3, cm^3, or dm^3
• Density: mass per unit volume, SI unit is kg/m^3, often expressed in g/cm^3

Temperature

• Measured in Celsius (°C), Fahrenheit (°F), or Kelvin (K) scales
• Conversion between Celsius and Kelvin scales: K = °C + 273.15

Scientific Notation

• Representation of numbers in the form N × 10^n (where n is an exponent and N is between 1 and 10)
• Example: 232.508 can be written as 2.32508 × 10^2 in scientific notation

Significant Figures and Rounding Off

• Significant figures: indicate the reliability of a measurement, include all certain digits and one doubtful digit
• Rules for determining significant figures:
  • Non-zero digits are significant
  • Zeros between non-zero digits are significant
  • Leading zeros are not significant
  • Trailing zeros are significant if there is a decimal point
  • In exponential notation, the numerical portion represents the number of significant figures
• Rounding off: procedure to retain the required number of significant figures

Dimensional Analysis

• Conversion of units from one system to another
• Factor label method or unit factor method

Physical Classification of Matter

• Solid: definite volume and shape, high intermolecular force of attraction, small intermolecular space, not compressible, rigid, and does not flow
• Liquid: definite volume, indefinite shape, moderate intermolecular force of attraction, slightly greater intermolecular space, not compressible, and flows
• Gas: indefinite volume and shape, negligible intermolecular force of attraction, large intermolecular space, compressible, and flows

Chemical Classification of Matter

• Elements: simplest form of matter, cannot be split into simpler substances, classified into metals, non-metals, and metalloids
• Compounds: pure substances made up of two or more elements combined in a definite proportion by mass
• Mixtures: combinations of two or more elements or compounds in any proportion, classified into homogeneous and heterogeneous mixtures

Laws of Chemical Combinations

• Law of Conservation of Mass: matter cannot be created or destroyed
• Law of Definite Proportions: chemical compounds always consist of the same elements combined together in the same ratio
• Law of Multiple Proportions: masses of one element that combine with a fixed mass of another element bear a simple ratio to one another
• Gay Lussac's Law of Gaseous Volumes: gases combine or are produced in a chemical reaction in a simple ratio by volume
• Avogadro Law: equal volumes of gases at the same temperature and pressure contain equal number of molecules

Atoms and Molecules

• Atom: smallest unit of an element, indivisible, and has mass
• Molecule: smallest unit of a substance, capable of independent existence, classified into homoatomic and heteroatomic
• Atomic Mass Unit (amu): defined as one twelfth the mass of a carbon-12 atom
• Atomic Mass: average relative mass of an atom of an element
• Molecular Mass: average relative mass of a molecule of a substance

Mole Concept

• Mole: amount of a substance that contains the same number of chemical units as there are atoms in exactly 12 grams of pure carbon-12
• Mole represents a collection of 6.022 × 10^23 chemical units
• Molar Mass: mass of one mole of a substance in grams
• Molar Volume: volume occupied by one mole of any substance, equal to 22.4 liters or 22,400 mL at 273 K and 1 atm pressure### Types of Mixtures
• Homogeneous mixture: A mixture that has the same composition throughout the sample.
• Heterogeneous mixture: A mixture that consists of two or more phases, which have different compositions and can be seen with the naked eye.

Law of Definite Proportions

• States that a chemical compound always consists of the same elements combined together in the same ratio, regardless of the method of preparation or source.

Empirical Formula

• The simplest whole number ratio of atoms of different elements present in a compound.
• Examples: N2O4 (NO2), C6H12O6 (CH2O), H2O (H2O), H2O2 (HO)

Precision and Accuracy

• Precision: The closeness of various measurements for the same quantity.
• Accuracy: The agreement of a particular value to the true value of the result.

Law of Multiple Proportions

• States that when two elements combine to form two or more compounds, the different masses of one element that combine with a fixed mass of the other bear a simple ratio to one another.
• Example: Carbon combines with oxygen to form CO and CO2, with a simple ratio of 16:32 or 1:2.

Atomic Mass

• The average relative mass of an atom of an element compared to the mass of an atom of carbon-12 taken as 12.
• Example: Chlorine has two isotopes of atomic mass units 34.97 and 36.97, with a relative abundance of 0.755 and 0.245 respectively, resulting in an average atomic mass of 35.46 u.

Percentage Composition

• The percentage of each element in a compound by mass.
• Example: Water (H2O) has a percentage composition of 11.18% hydrogen and 88.79% oxygen.

Significant Figures

• Rules for rounding off:
  1. If the digit coming after the desired number of significant figures is more than 5, the preceding significant figure is increased by one.
  2. If the digit involved is less than 5, it is neglected and the preceding significant figure remains unchanged.
  3. If the digit happens to be 5, the last mentioned or preceding significant figure is increased by one only in case it happens to be odd, and it remains unchanged in case of an even figure.

Unit Factor Method

• A method to convert units from one system to another.
• Example: Converting 93 million miles to SI units.

Compound vs Mixture

• Key differences:
  • Constituents are always present in a fixed ratio by mass (compound) vs may be present in any ratio (mixture).
  • May or may not be homogeneous in nature (compound) vs always homogeneous (mixture).
  • Constituents can be easily separated by simple mechanical means (mixture) vs cannot be easily separated (compound).

Molarity

• The number of moles of solute dissolved per liter (dm3) of the solution.
• Example: Calculating the molarity of NaOH in a solution prepared by dissolving 4g of NaOH in enough water to form 250mL of the solution.

Classification of Substances

• Pure substances: Ethyl alcohol, oxygen, carbon, distilled water.
• Mixtures: Blood, steel.

Formula Mass

• The sum of atomic masses of the elements present in a formula unit of a compound.
• Example: Calculating the formula mass of NaOH.

Scientific Notation

• Expressing numbers in the form of a x 10^n, where a is a number between 1 and 10, and n is an integer.
• Examples: 0.0048 (4.8 x 10^-3), 234,000 (2.3 x 10^5), 200.0 (2.0 x 10^2).

Test your knowledge of fundamental chemistry concepts, including branches of chemistry, properties of matter, and SI units.