Chim Chum Chem

• Displacement of water:

  • Method for collecting gases which are insoluble or slightly soluble in water
  • Gases rise to the top of the gas jar if they do not dissolve in water
  • Examples of gases collected via this method: H2, O2, CO, and CO2
• Upwards delivery:

  • Method used to collect gases with a lighter density compared to air
  • Gases like NH3 & H2 can be collected using this method
  • The molar mass of air is around 28.8 (78% N2 + 21% O2)
• Downwards delivery:

  • Method used to collect gases with a heavier density compared to air
  • Examples of gases collected using this method: Cl2, HCl, and SO2
  • The molar mass of air is around 28.8 (78% N2 + 21% O2)
• Drying of gas:

  • To dry a sample of gas, pass it through drying agents like concentrated sulfuric acid, quicklime (calcium oxide), and fused calcium chloride
• There are 9 different separation techniques, including filtration, evaporation, crystallisation, sublimation, and magnetic separation
• Pure substances contain only one type of substance, while impure mixtures contain two or more substances
• Compounds are chemically combined, while mixtures are physically combined
• Compounds can be separated using chemical methods, while mixtures can be separated using physical methods (separation techniques)
• Simple distillation involves heating the solution in a round-bottomed flask, while fractional distillation separates liquids based on their boiling points using a fractionating column
• Fractional distillation requires a minimum of at least 10°C difference in boiling points between the miscible liquids in the solution
• Factors affecting the differences in Rf values in chromatography include the solubility of substances in the solvent used and the molecular masses of the substances
• Test for Cations:

  • Cation reaction with NaOH (aq) (strong alkaline) and NH3 (aq) (weak alkaline)
  • Copper(II) Cu2+: Forms blue precipitate of Cu(OH)2, which is insoluble in excess NaOH and dissolves in excess NH3 to give dark blue complex ion
  • Iron(II) Fe2+: Forms green precipitate of Fe(OH)2, which is insoluble in excess NaOH and NH3
  • Iron(III) Fe3+: Forms reddish-brown precipitate of Fe(OH)3, which is insoluble in excess NaOH and NH3
  • Calcium Ca2+: Forms white precipitate of Ca(OH)2, which is insoluble in excess NaOH
• Basics to know about the states of matter:

  • Solid: fixed volume and shape, cannot be compressed, does not flow
  • Liquid: fixed volume, no fixed shape, cannot be compressed, flows easily
  • Gas: no fixed volume or shape, can be compressed easily, flows in all directions
• Kinetic Particle Theory of Matter:

  • All matter consists of particles that are too small to be directly visible
  • Particles are always in a constant state of random motion at varying speeds
• Physical properties based on the particulate model of matter:

  • Solid: closely packed in an orderly arrangement, very strong attractive force, very high density, particles vibrate about their fixed position
  • Liquid: loosely packed in a disorderly arrangement, strong attractive force, high density, particles slide over one another freely
  • Gas: far apart and random arrangement, weak attractive force, low density, particles move about at high speeds randomly
• Change in state:

  • Melting (Solid to Liquid):
  • Particles gain energy from surroundings and vibrate vigorously about their fixed positions
  • Temperature remains constant during melting process
  • Mixture of solid and liquid present
  • Freezing (Liquid to Solid):
  • Particles lose kinetic energy and freezing starts
  • Temperature remains constant as heat energy is released to surroundings
  • Mixture of solid and liquid present
  • Boiling (Liquid to Gas):
  • Heat energy absorbed reaches boiling point
• • Particles gain energy to overcome forces of attraction and move further apart
  • Temperature remains constant during boiling process
  • Mixture of gas and liquid present
  • Condensation (Gas to Liquid):
  • Particles lose kinetic energy and condensation starts
  • Heat energy is released to surroundings
  • Temperature remains constant
  • Mixture of gas and liquid present
• Sublimation & Deposition:

  • Sublimation: solid to gas transition (e.g., iodine)
  • Deposition: gas to solid transition (e.g., dry ice)
• Diffusion:

  • Movement of molecules from a region of higher concentration to a region of lower concentration
  • Gas or liquid particles move to available spaces through random motion
  • Factors affecting diffusion:
  • Heavier molecules move slower
  • Higher temperature leads to faster diffusion
  • Lower mass particles result in faster diffusion
  • Gaseous state has faster diffusion than liquid state
  • Greater concentration gradient leads to faster diffusion
• Subatomic particles:

  • Proton:
  • Charge: +1
  • Relative mass: 1
  • Symbol: p
  • Location: Nucleus
  • Neutron:
  • Charge: 0
  • Relative mass: 1
  • Symbol: n
  • Location: Nucleus
  • Electron:
  • Charge: -1
  • Relative mass: 1/1836 (negligible mass)
  • Symbol: e
  • Location: Electron shell
• Electronic configurations:

  • Must know: 2,8,8 electronic configuration
  • For elements after calcium, the third shell can hold a maximum of 18 electrons
  • First shell: Maximum of 2 electrons
  • Second shell: Maximum of 8 electrons
  • Third shell: Maximum of 8 electrons
  • Ar atom has 18 protons and 22 neutrons
• Proton number:

  • Total number of protons in an atom (number of electrons as well)

  Nucleon number:

  • Total number of protons and neutrons in the nucleus of an atom
  • Identity of an element depends on its proton number, not its nucleon number
  • To review later: Isotopes
• Formation of positive ions:

  • Atoms that lose electrons become positively charged
  • More protons than electrons
  • Become cations
  • Example: Sodium atom loses one electron to become a sodium cation with a charge of +1 (Na+)
• Formation of negative ions:

  • Atoms that gain electrons become negatively charged
  • More electrons than protons
  • Example: Chlorine atom gains one electron to become a chlorine anion with a charge of -1 (Cl-)
• Isotopes:

  • Atoms of the same element with the same amount of protons and electrons but different amounts of neutrons
  • Same chemical properties as they have the same amount of electrons
  • Differences in physical properties due to different neutron amounts affecting mass and other physical properties like density
  • Example: Chlorine exists as chlorine-35 and chlorine-37 atoms, with different percentage abundances contributing to the average atomic mass of 35.5
• Things to note about isotopes:

  • Different number of neutrons causes differences in physical properties such as density
  • Same number of protons/electrons
  • Similar chemical properties as atoms undergo the same chemical reactions to form compounds with the same chemical formula
  • Atomic mass is an average mass of the element's isotopes, calculated based on percentage composition and respective masses
  • Example: Chlorine's Ar is 35.5
• Elements, compounds, and mixtures:

  • Elements are naturally found and can exist by themselves or in diatomic molecule form for gases such as H2 or O2
  • Compounds are combined using chemical methods, have a fixed ratio, and fixed melting and boiling points
  • Mixtures are combined using physical methods, can have any ratio, and melt and boil over a range of temperatures
• Definition of a molecule:

  • A molecule is when 2 or more atoms chemically combine
  • An element can also exist as a diatomic molecule
  • A compound is defined as 2 or more elements chemically combined, hence a compound must be a molecule
• 3 types of bonds:

  • Ionic, covalent, metallic
  • Atoms form chemical bonds to achieve a stable electronic configuration (2,8,8) by transferring electrons, sharing electrons, or forming a metal lattice
  • Non-metal atoms can form ionic bonds with metal atoms and covalent bonds with other non-metal atoms
  • Metal atoms can form metallic bonds with other metal atoms
• Ionic bonds:

  • Formed between metals and non-metals
  • Transfer of electrons from metal to non-metal for both to have complete valence shells
  • Result in oppositely charged ions attracting each other
  • Ionic compounds have a giant ionic lattice structure held by strong electrostatic forces of attraction
  • Physical properties include high melting and boiling points, solubility in water, ability to conduct electricity in molten and aqueous states, poor heat conduction, and not being volatile
• Covalent bonds:

  • Formed between non-metal atoms
  • Defined as the electrostatic force of attraction between nuclei of atoms and shared electrons
  • Atoms share valence electrons to attain stable electronic configurations
  • Can be simple molecular or giant molecular structures
  • Simple molecular structures have low melting and boiling points due to weak intermolecular forces of attraction, while giant molecular structures have high melting and boiling points due to strong covalent bonds
• Simple molecular structure:

  • Physical properties include low melting and boiling points, insolubility in water, inability to conduct electricity in any state, poor heat conduction, and high volatility
  • Melting or boiling only requires breaking weak intermolecular forces between molecules, not breaking covalent bonds within the molecule itself
  • Simple molecular substances are usually gases or liquids at room temperature and are unable to conduct electricity
• Simple molecular substances are insoluble in water
• Simple molecular substances are unable to conduct electricity due to the absence of mobile charged carriers (electrons or ions)
• Simple molecular substances are poor conductors of heat
• Giant molecular structures are held together by strong covalent bonds
• Giant molecular structures have high melting and boiling points due to the strong covalent bonds that require a huge amount of energy to overcome
• Examples of giant molecular structures include Diamond, Graphite, and Silicon Dioxide (Sand)
• Diamond is comprised of carbon atoms held together by strong single covalent bonds
• Diamond has a tetrahedral structure and is known for its hardness