Basic Concepts of Chemistry

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Questions and Answers

Which of the following statements about metalloids is accurate?

  • Metalloids are excellent conductors of heat and electricity.
  • Metalloids include metals like iron and copper.
  • Metalloids have no applications in technology.
  • Metalloids possess properties that are intermediate between metals and non-metals. (correct)

What defines the physical properties of matter?

  • They cause a change in the chemical composition of the substance.
  • They depend on the state of matter being solid only.
  • They can be measured without altering the chemical structure. (correct)
  • They can only be observed during a chemical reaction.

Which of the following is NOT a base physical quantity?

  • Color (correct)
  • Electric Current
  • Mass
  • Thermodynamic Temperature

How do the particles in a gas compare to those in a solid?

<p>Gas particles are far apart, unlike the organized structure of solid particles. (B)</p> Signup and view all the answers

What distinguishes mass from weight?

<p>Mass remains constant regardless of location, while weight can vary. (A)</p> Signup and view all the answers

Which of the following statements accurately reflects the characteristics of pure substances?

<p>They have a definite chemical composition and consistent properties. (D)</p> Signup and view all the answers

Which branch of chemistry focuses specifically on the study of carbon compounds?

<p>Organic chemistry (A)</p> Signup and view all the answers

What distinguishes a mixture from a pure substance?

<p>Mixtures do not have a definite chemical composition and therefore lack consistent properties. (A)</p> Signup and view all the answers

What are the classifications of elements in chemistry based on their properties?

<p>Metals, nonmetals, and metalloids (D)</p> Signup and view all the answers

What type of chemistry primarily involves the identification and quantification of chemical substances?

<p>Analytical chemistry (B)</p> Signup and view all the answers

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Study Notes

Introduction to Chemistry

  • Chemistry involves the study of matter, its properties, and the changes it undergoes.
  • Matter is divided into pure substances and mixtures based on chemical composition.

Pure Substances vs Mixtures

  • Pure substances: have a fixed composition and consistent properties; examples include distilled water and pure metals.
  • Mixtures: do not have fixed compositions; examples include paint and concrete.

Branches of Chemistry

  • Organic chemistry: focuses on carbon compounds.
  • Inorganic chemistry: deals with substances not covered by organic chemistry.
  • Physical chemistry: studies the principles and physical properties of matter.
  • Biological chemistry: examines the chemistry of living organisms.
  • Analytical chemistry: involves the identification and quantification of chemical substances.

Classification of Pure Substances

  • Pure substances are categorized into:
    • Elements: Cannot be broken down into simpler substances; classified as metals, nonmetals, and metalloids.
    • Compounds: Combinations of two or more elements.

Properties of Matter

  • Metals: Lustrous, conductive, ductile, malleable; examples include gold and copper.
  • Non-metals: Dull, poor conductors, brittle; examples include nitrogen and iodine.
  • Metalloids: Have mixed characteristics; examples include silicon and arsenic.

States of Matter

  • Solid: Particles tightly packed in a structured arrangement.
  • Liquid: Particles close together but can move.
  • Gas: Particles are widely spaced.

Properties of Matter

  • Physical properties: Can be observed without changing composition (e.g., color, odor, melting point).
  • Chemical properties: Only observable during chemical reactions (e.g., flammability, reactivity).

Measurement of Properties

  • Mass and Weight: Mass is constant and does not change, while weight varies due to gravity. SI unit for mass is the kilogram (kg).
  • Length: Measured in meters (m); small lengths can be expressed in nanometers (nm) and picometers (pm).
  • Volume: Amount of space occupied, measured in cubic meters (m³) or liters (L).

SI Units

  • The International System of Units (SI) defines seven base units:
    • Length: metre (m)
    • Mass: kilogram (kg)
    • Time: second (s)
    • Electric Current: ampere (A)
    • Temperature: kelvin (K)
    • Amount of Substance: mole (mol)
    • Luminous Intensity: candela (cd)

Laws of Chemical Combination

  • Law of Conservation of Mass: Mass cannot be created or destroyed (Lavoisier).
  • Law of Definite Proportions: Compounds have a consistent ratio of elements by mass (Proust).
  • Law of Multiple Proportions: Different compounds formed from the same elements will have mass ratios of small whole numbers (Dalton).
  • Gay-Lussac’s Law: Gases react in simple volume ratios at constant temperature and pressure.

Atomic Theory and Atomic Mass

  • Dalton's Atomic Theory: Matter consists of indivisible atoms, and atoms combine in fixed ratios to form compounds.
  • Atomic mass: Measured in atomic mass units (amu); 1 amu is one-twelfth the mass of a carbon-12 atom.

Average Atomic Mass

  • Average atomic mass is the weighted average based on isotopes and their abundances.
  • Example for carbon: average atomic mass is 12.011 u based on isotopes (^{12}C), (^{13}C), and (^{14}C).

Moles and Molar Mass

  • A mole contains Avogadro's number (6.022 \times 10^{23}) particles.
  • Molar mass: Mass in grams per mole corresponds to the atomic mass in amu.

Molecular Mass and Formula Mass

  • Molecular mass: Total average atomic mass of all atoms in a molecule.
  • Formula mass: Sum of atomic masses in ionic compounds like NaCl, where discrete molecules do not exist.

Calculations Examples

  • Molecular mass of water (Hâ‚‚O) = 2(1.0 u) + 16.0 u = 18.0 u.
  • Formula mass of NaCl = 23.0 u (Na) + 35.5 u (Cl) = 58.5 u.### Atomic and Molar Mass
  • Atomic mass is expressed in atomic mass units (u), while molar mass is the mass of one mole of a substance in grams per mole (g mol-1).
  • Examples of elemental atomic and molar masses include:
    • Hydrogen (H): 1.0 u, 1.0 g mol-1
    • Carbon (C): 12.0 u, 12.0 g mol-1
    • Oxygen (O): 16.0 u, 16.0 g mol-1
  • For polyatomic substances, molar mass numerically equals molecular/formula mass in u.

Polyatomic Substances

  • Common polyatomic substances and their respective molar and molecular masses:
    • Oxygen (O2): 32.0 u, 32.0 g mol-1
    • Water (H2O): 18.0 u, 18.0 g mol-1
    • Sodium Chloride (NaCl): 58.5 u, 58.5 g mol-1

Urea Example Calculation

  • Urea's molecular formula: NH2CONH2
  • Calculation of molecular mass:
    • 2 Nitrogen (N): 2 x 14 u
    • 1 Carbon (C): 1 x 12 u
    • 4 Hydrogen (H): 4 x 1 u
    • 1 Oxygen (O): 1 x 16 u
  • Total molecular mass = 60 u
  • Molar mass of urea = 60 g mol-1
  • Number of moles in 5.6 g of urea:
    • Moles = mass / molar mass = 5.6 g / 60 g mol-1 = 0.0933 mol
  • Number of molecules using Avogadro's constant:
    • Molecules = 0.0933 mol x 6.022 x 10^23 molecules/mol ≈ 5.618 x 10^22 molecules

Atom Calculation Examples

  • For 52 moles of Argon (Ar):
    • Atoms = 52 moles x 6.022 x 10^23 atoms/mol ≈ 313.144 x 10^23 atoms of Ar
  • For 52 u of Helium (He):
    • Atomic mass of He = 4.0 u, thus 52 u corresponds to:
    • Atoms = 52 / 4.0 ≈ 13 atoms of He
  • For 52 g of Helium:
    • Molar mass of He = 4.0 g mol-1
    • Moles = 52 g / 4.0 g mol-1 = 13 mol
    • Atoms = 13 mol x 6.022 x 10^23 atoms/mol ≈ 78.286 x 10^23 atoms of He

Ammonia Gas Calculation

  • Volume of ammonia (NH3) at STP = 67.2 dm3
  • Molar volume of a gas at STP = 22.4 dm3/mol
  • Number of moles of NH3:
    • Moles = Volume / Molar volume = 67.2 dm3 / 22.4 dm3/mol = 3.0 mol
  • Number of molecules in 3.0 moles:
    • Molecules = 3.0 mol x 6.022 x 10^23 molecules/mol ≈ 18.066 x 10^23 molecules

Understanding Moles and Gases

  • Moles of gases can be calculated using volume rather than mass.
  • Avogadro's law indicates that one mole of any gas occupies 22.4 dm3 at standard temperature (0°C) and pressure (1 atm), known as the molar volume of a gas.

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