IBDP Chemistry - Structure & Reactivity Notes

Structure & Reactivity

• Two main concepts in IBDP Chemistry:
  • Structure
  • Reactivity
• Structure determines reactivity, which in turn transforms structure.

Structure

• One of the two main concepts in the DP chemistry course.
• Concerned with the way in which the particles of a substance are arranged.
• Importance: The structure of different substances can affect their physical properties (e.g., melting and boiling points) and chemical reactivity.
• Structure topics:
  • S1: Models of the particulate nature of matter
  • S2: Models of bonding & structure
  • S3: Classification of matter

Reactivity

• A chemical reaction involves the breaking of bonds and the formation of new bonds.
  • Breaking bonds requires energy.
  • Energy is released when new bonds are formed.
  • The change in energy is the driving force behind reactions.
• Reactivity: Concerned with how and why substances interact with each other.
• Reactivity topics:
  • R1: What drives chemical reactions
  • R2: How much, how fast, how far?
  • R3: What are the mechanisms of chemical change?

Review of MYP Science

• Background knowledge review through worksheets.
• Topics to be confident in:
  • Naming ions
  • Naming ionic compounds
  • Naming polyatomic ions
  • Naming covalent compounds (less important, as most compounds used are ionic).

Structures 1.1: Introduction to the Particulate Nature of Matter

• Guiding Question: How is the particulate nature of matter modelled?

Learning Outcomes

• Understand:
  • Elements are the primary constituents of matter and cannot be chemically broken down into simpler substances.
  • Compounds consist of atoms of different elements chemically bonded together in a fixed ratio.
  • Mixtures contain more than one element or compound in no fixed ratio, which are not chemically bonded and can be separated by physical methods.
  • The Kinetic Molecular Theory (KMT) is a model to explain physical properties of matter (solids, liquids, and gases) and changes of state.
  • Temperature (in Kelvin) is a measure of average kinetic energy of particles.
• Apply Knowledge to:
  • Distinguish between the properties of elements, compounds, and mixtures.
  • Convert between values in the Celsius and Kelvin scales.
  • Use state symbols (s, l, g, and aq) in chemical equations.
  • Distinguish between the different states of matter.
  • Interpret observable changes in physical properties and temperature during changes of state.

Links to Other Topics

• Factors considered in choosing a method to separate mixture components.
• How products of a reaction can be purified.
• Influence of intermolecular forces on the type of mixture formed between two substances (S2.2).
• Why alloys are generally considered mixtures, even with metallic bonding (S2.3).
• Reasons for some substances being solid while others are fluid under standard conditions (S2.4).
• Why some changes of state are endothermic and some exothermic (S2, R1.2).
• Graphical distribution of kinetic energy values of particles in a sample at a fixed temperature (R2.2).
• What must happen to particles for a chemical reaction to occur (R2.2).

Clarification Notes

• Understanding of solvation, filtration, evaporation, distillation, reflux, and paper chromatography is required.
• Differences between homogeneous and heterogeneous mixtures should be understood.
• Names of changes of state (melting, freezing, vaporization (evaporation and boiling), condensation, sublimation, and deposition) should be understood.
• The kelvin (K) is the SI unit of temperature and has the same incremental value as the Celsius degree (°C).

Theory of Knowledge/Nature of Science

• In the early nineteenth century John Dalton (1766–1844) set out values for atomic weights (now atomic masses) of the elements.
• He demonstrated that two elements (A and B) combine to make a compound in one or more set ratios, AB, AB2, A2B, etc.
• This paved the way for chemical formulas and chemical equations.
• Importance of obtaining compounds in a pure state and making reliable and reproducible quantitative measurements.

Real World Context

• Chromatography is a useful way to obtain pure substances from mixtures containing very small amounts of the individual components.
• Coupled with techniques like mass spectrometry, it can separate and identify complex mixtures both quantitatively and qualitatively.
• Different types of chromatography: paper, thin layer (TLC), column (LC), gas-liquid (GLC), and high-performance liquid chromatography (HPLC).
• Each type involves a stationary phase and a mobile phase.
• Chromatography relies on components having different tendencies to adsorb onto a surface or dissolve in a solvent.
• Adsorption involves a solid stationary phase with a moving liquid phase (LC, HPLC, at times TLC).
• The rate at which the solute moves through the solid phase depends on its solubility in the mobile phase and its degree of adsorption onto the stationary phase.

Elements & Compounds

• All pure substances are either made up of just one chemical element or are compounds made up of two or more elements.
• A chemical element is a neutral substance that cannot be broken down into simpler substances using chemical methods.
  • Consists only of atoms that contain the same number of protons, though atoms can contain a different number of neutrons.
  • There are 118 known chemical elements in the periodic table.
• Some elements exist as diatomic molecules (e.g., H2, Cl2, O2, N2).
• A molecule is an electrically neutral group of two or more atoms bonded together.
• A compound is a pure substance formed when two or more chemical elements are chemically bonded together.
• A compound has different properties from its component elements.
  • Example: Sodium is a reactive metal (melts at 98 °C), and chlorine is a poisonous reactive gas, but sodium chloride is a salt with a high melting point (801 °C) that dissolves in water.

Particles

• Particles are tiny pieces of matter.
• Atoms are the smallest pieces of an element.
  • Example: The smallest piece of hydrogen is 1 atom of hydrogen.
• Molecules consist of 2 or more atoms bonded together.
  • Examples: \text{H}2[0.1cm]\text{O}, \text{O}2, NaCl, \text{H}_2.
• Compounds consist of 2 or more different atoms bonded together.
  • Examples: \text{H}_2O and NaCl.

States of Matter: Properties of Solids, Liquids, and Gases

Heat

• Heat is a measure of the total amount of energy in a given amount of substance and depends upon the amount of substance present.
• Measured in joules.
• Heat is a form of energy that always transfers from a higher temperature object to a lower temperature object.

Temperature

• Temperature is a measure of the 'hotness' of a substance and is independent of the amount of substance present.
• It is a measure of the average kinetic energy of the substance.
• Absolute temperature is measured in kelvin.
• 0 K (‘zero Kelvin’, not ‘okay’!), or “absolute zero”, is the temperature at which all motion has stopped.
• Temperature cannot go below this value.
• A change in 1 K is the same as a change in 1 °C.

Changes of State - Kinetic Theory of Matter

• Solid State: Particles are held closely together in fixed positions in a lattice. Particles can vibrate about a fixed point but possess no individual translational velocity.
• Liquid State: As heat is supplied, the vibration intensifies, and eventually, the particles have sufficient energy for the lattice to break, and the particles are free to move. There are still attractive forces between the particles.
• Gaseous State: As more heat is added, the particles move faster, i.e., gain kinetic energy. Eventually, they will have sufficient energy to overcome the attractive forces and escape as a vapor. When the vapor pressure equals the external pressure, the liquid boils. Hence, the boiling point depends upon the external pressure. At 100 kPa pressure, water boils at 373 K (100 °C).

• Changes of state terms:
  • Melting
  • Freezing
  • Evaporation
  • Condensation
  • Sublimation
  • Deposition
• Note the differences between boiling, vaporization, and evaporation

Pure Substances and Mixtures

• A pure substance may be an element or a compound and has a definite and constant composition.
• A mixture is a physical combination of pure substances. Mixtures have no definite or constant composition.
• A mixture can be separated into its pure components by physical means.

Separating Mixtures

• Ways of separating mixtures into their pure components include: solvation, filtration, evaporation, distillation (also fractional distillation), and different forms of chromatography (paper chromatography, thin-layer chromatography, and gas-liquid chromatography).
• The choice of separation technique depends on what is in the mixture, the physical and chemical properties of the substances present, and whether the substances to be separated are solids, liquids, or gases.
• Homogeneous mixture: all the components are in the same phase
• Heterogeneous mixture: the components are in different phases. There is a physical boundary between two phases.
• Separation Techniques:
  1. Filtration: Separates solid particles from a liquid. Can be carried out under gravity or reduced pressure.
  2. Solvation and evaporation: Warming a mixture with a solvent to dissolve a soluble component, filtering to remove insoluble particles, then evaporating the solvent.
  3. Recrystallization: Dissolving an impure solid in a suitable solvent by heating, cooling slowly to form crystals, washing the crystals, and drying.
  4. Distillation: Used to separate a volatile liquid from dissolved components or less volatile liquids.
  5. Reflux: Used to heat volatile mixtures without losing any of the volatile components.
  6. Chromatography: Used to separate different components in a mixture. In paper chromatography a spot of the mixture is placed on the paper and the solvent (eluent) allowed to rise up the paper. Different components partition between the solvent and the paper by different amounts so rise at different rates.
• R_f = distance travelled by component / distance travelled by the solvent from the original spot.

John Dalton

• In the early nineteenth century John Dalton made the following assumptions.
  1. Elements are made of extremely small particles called atoms.
  2. Atoms of a given element are identical in size, mass, and other properties; atoms of different elements differ in size, mass, and other properties.
  3. Atoms cannot be subdivided, created, or destroyed.
  4. Atoms of different elements combine in simple whole-number ratios to form chemical compounds.
  5. In chemical reactions, atoms are combined, separated, or rearranged.