Eton College Chemistry: Atoms, States of Matter, and Rates of Reaction Notes

Atoms, Molecules, and Atomic Structure

• Definition of an Atom: An atom represents the smallest part of an element that can exist. It is a fundamental building block of matter.

• Definition of a Molecule: A molecule is a group of two or more atoms chemically bonded together. These can be atoms of the same element (e.g., $O_2$) or different elements (e.g., $H_2O$).

• Subatomic Particles and Atomic Structure:

  • Proton: Found in the nucleus. It has a relative mass of $1$ and a relative charge of $+1$.

  • Neutron: Found in the nucleus. It has a relative mass of $1$ and a relative charge of $0$.

  • Electron: Found in the shells (energy levels) surrounding the nucleus. It has a relative mass of $\frac{1}{1840}$ (effectively negligible) and a relative charge of $-1$.

• Fundamental Atomic Terms:

  • Atomic Number: The number of protons found in the nucleus of an atom. This number is unique to each element and determines its position in the Periodic Table.

  • Mass Number: The total sum of protons and neutrons in the nucleus of an atom.

  • Isotopes: Atoms of the same element that contain the same number of protons but a different number of neutrons. Consequently, isotopes have the same atomic number but different mass numbers.

• The Periodic Table:

  • Elements are arranged in order of increasing atomic number.

  • Groups: Vertical columns that contain elements with similar chemical properties (due to having the same number of electrons in their outer shell).

  • Periods: Horizontal rows that represent the number of electron shells in the atoms.

• Required Element Names and Symbols:

  • Group 1 & 2: Hydrogen ($H$), Helium ($He$), Lithium ($Li$), Beryllium ($Be$), Sodium ($Na$), Magnesium ($Mg$), Potassium ($K$), Calcium ($Ca$), Rubidium ($Rb$), Strontium ($Sr$), Caesium ($Cs$), Barium ($Ba$), Francium ($Fr$), Radium ($Ra$).

  • Group 3-8/0: Boron ($B$), Carbon ($C$), Nitrogen ($N$), Oxygen ($O$), Fluorine ($F$), Neon ($Ne$), Aluminium ($Al$), Silicon ($Si$), Phosphorus ($P$), Sulfur ($S$), Chlorine ($Cl$), Argon ($Ar$).

  • Halogens & Others: Bromine ($Br$), Iodine ($I$), Astatine ($At$), Tin ($Sn$), Lead ($Pb$).

  • Transition Metals: Iron ($Fe$), Nickel ($Ni$), Copper ($Cu$), Zinc ($Zn$), Palladium ($Pd$), Silver ($Ag$), Platinum ($Pt$), Gold ($Au$), Mercury ($Hg$).

Particle Motion and States of Matter

• Evidence for Particles:

  • Dilution of Coloured Solutions: When a crystal of potassium manganate(VII) is dissolved in water, the purple color spreads. Further dilution makes the color lighter but still visible, demonstrating that the matter is composed of many tiny particles that redistribute themselves through the solvent.

  • Diffusion of Gases: The movement of particles from an area of high concentration to an area of low concentration. A classic example is the reaction between ammonia ($NH_3$) and hydrogen chloride ($HCl$) gases to form a white ring of ammonium chloride ($NH_4Cl$). The ring forms closer to the concentrated hydrochloric acid because ammonia particles are lighter and diffuse faster.

• The Three States of Matter:

  • Solid: Particles are arranged in a regular, fixed lattice. They vibrate about fixed positions and have the lowest kinetic energy.

  • Liquid: Particles are arranged randomly but remain in contact (touching). They can move around each other and have more kinetic energy than solids.

  • Gas: Particles are arranged randomly and are far apart. They move rapidly in all directions and have the highest kinetic energy.

• Interconversions of State:

  • Melting: Solid to liquid. Achieved by heating until particles gain enough energy to overcome the forces holding them in the lattice.

  • Freezing: Liquid to solid. Achieved by cooling until particles lose energy and become fixed in a lattice.

  • Evaporation/Boiling: Liquid to gas. Achieved by heating until particles gain enough kinetic energy to break all attractive forces between them.

  • Condensation: Gas to liquid. Achieved by cooling; particles lose energy and move closer together.

  • Sublimation: Solid to gas directly (e.g., dry ice or iodine).

  • Deposition: Gas to solid directly.

Classification and Separation of Substances

• Definitions:

  • Element: A substance consisting of only one type of atom.

  • Compound: A substance made of two or more different elements chemically bonded in fixed proportions.

  • Mixture: A physical combination of two or more substances that are not chemically joined and can be separated by physical means.

• Purity and Thermal Behavior:

  • Pure Substances: Characterized by a fixed, sharp melting point and boiling point (e.g., pure water boils at exactly $100\,^{\circ}\text{C}$).

  • Mixtures: These do not have a fixed melting or boiling point; instead, they melt or boil over a range of temperatures.

• Separation Techniques:

  • Simple Distillation: Used to separate a solvent from a solution (e.g., obtaining pure water from salt water).

  • Fractional Distillation: Used to separate a mixture of miscible liquids with different boiling points (e.g., ethanol and water).

  • Filtration: Used to separate an insoluble solid from a liquid.

  • Crystallisation: Used to separate a soluble solid from a solution by evaporating the solvent until it reaches saturation, then allowing crystals to form.

  • Paper Chromatography: Used to separate substances based on their different solubilities in a specific solvent.

• Chromatography and Rf Values:

  • A chromatogram shows the components of a mixture as separate spots.

  • $R_f$ Calculation: The Retention Factor ($R_f$) is calculated as:

$R_f = \frac{\text{distance moved by the substance}}{\text{distance moved by the solvent front}}$

• The $R_f$ value is always between $0$ and $1$. It allows scientists to identify components by comparing values against known standards.

Rates of Reaction

• Factors Affecting Reaction Rate:

  • Surface Area of a Solid: Increasing the surface area (e.g., using powder instead of marble chips) increases the rate.

  • Concentration of a Solution: Increasing the number of particles per unit volume increases the rate.

  • Pressure of a Gas: Increasing pressure pushes gas particles closer together, increasing the rate.

  • Temperature: Increasing temperature increases the kinetic energy of particles, leading to a higher rate.

  • Use of a Catalyst: Speeds up the reaction without being consumed.

• Collision Theory Explanations:

  • For a reaction to occur, particles must collide with sufficient energy (activation energy) and in the correct orientation.

  • Surface Area: More particles are exposed at the surface, leading to more frequent successful collisions.

  • Concentration/Pressure: Particles are more crowded, leading to a higher frequency of collisions per unit of time.

  • Temperature: Particles move faster, resulting in more frequent collisions. Crucially, a much higher proportion of particles possess energy greater than the activation energy ($E_a$).

• Catalysts:

  • Definition: A substance that increases the rate of a chemical reaction but remains chemically unchanged at the end of the process.

  • Mechanism: A catalyst works by providing an alternative reaction pathway that has a lower activation energy ($E_a$). This means more particles have enough energy to react upon collision.

• Core Practicals:

  • Marble Chips and Hydrochloric Acid: Investigating the reaction $CaCO_3(s) + 2HCl(aq) \rightarrow CaCl_2(aq) + H_2O(l) + CO_2(g)$. The rate is measured by the volume of $CO_2$ gas produced over time. Comparison of large chips vs. small chips (surface area) and varied concentrations of $HCl$.

  • Catalytic Decomposition of Hydrogen Peroxide: Investigating $2H_2O_2(aq) \rightarrow 2H_2O(l) + O_2(g)$. Different solids (e.g., Manganese(IV) oxide) are tested to observe their effect on the rate of oxygen production.